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TABLE OF SPECTRA. 




A a, R C D 

. , 2 ,° , , , 3 ,° ,, 4 ,° , 6 ,° , 6 i° i. 




y ! v y G H HI 

70 80 90 100 110 120 130 140 150 160. 170 

I I ... II. . . I .... I ... .11 .... I .... I .... I .... I . . u I I u I I .. . . | ill I I III I « M I I I M I II I I I I I I I I I I I I I I In I M ' I I I 


0w VV , | | 

. ... I ... . I .... II I . , . I l , I t I . . I I I t I I I i II r M l I l l M I (I I I M t I ( II ( I I I I t I I 


ill 111! 1111 



■Hill! ill! ffillll!- 


Ti 111 1 1 u 111 1 ill i ffi ill 1 1 11 1 f l°i 11 11 111 | 4 |°i 111 1 111 i 1 °i 111 1 1 r i rh 111 1 111 ' 11 1 11 n h 111 1 1M1 1 11 . 1 1 1111 1 1111 1 1 " 1 1 1 " 1 1 " " 1 111' I "I" I " " I " ' * I " r' I " " I " " I " " I " I 1 


?i, i. 1 11 M h i II1 1 ill fl°i ill u 111 |T n 11111 Th 111111 i.Ti II111 miT i-i iin i a 11 ; , I n i iT, i n 111111. Ml 11111111111 • 111 !* 11.1.1.. IM, 111 -lit 1111 |J •' I • 111 • ■ I" " I".' I" " I '1" I '"I 


o 10 20 30 40 50 60 70 80 <0 lOO 110 120 130 140 150 160 170 

|il ill mill ml in, I mi In ilm 1 11,1 1 ti 11 1 , n Ji i n 11 , i , 1 . m Li mJ m 1 1,,. 1 1111 1 11 nl 111 il 1111 1 11 n 1 1 m I m n 1 1111 1 1 m I nil 1 1 1111 1 1111 1 1 m I m 1 1 m 11 111111 ill m 11 n 11 


n 10 20 30 40 50 60 70 80 00 100 110 120 ( 1330 140 150 100 17C 

|| 11111! 11 11, n | I I ! j | 11 11 | I I | 1111 I i 111 11 I 11 I...111 H I I'll'l I HI 11 11111111111 I 11 111 ill I 11 11 111111 1111111 11 I 1111 III 11 11 1111 ill I I I M 1111 111111 M I I 111 1111111111 III 11 1 11111 


0 10 20 30 40 50 60 70 80 SO . 100 110 120 130 140 150 160 lj( 

1111111.| 1111 11111 I 1111 I I 1111 11 11 I I I I I 1111 I I 111 11 1111111 nil I III M 111 i 1 111 1 1 1 1 11 I 1 1 I I I I I I I I I i 11 I I I I I I 11 I I I I I I I I I I I I I 11 III I I I ill I I I i I I I I I I I 11 1111 111 111 I I 11111 11 ill 111111 11 


o 10 20 30 40 50 60 70 80 SO 100 110 120 130 140 150 160 170 

]1111'| 1111111111111111111111 11111 111111111111111 i ri 1111 m 11 n ■ I h ii I i ii lln’i 11 i i n 11111! 11111111111111111111111 m 1111111 i i.l i ill 111111111111111111111111M111111 ill 11111111 



0 10 20 30 40 50 60 70 80 SO 100 110 120 130 140 150 160 .170 

h 11111111111111111111111111 11111 1111111 ii 111111 11111111111111111111 mli 111111111111111111111111111111111 111111 m 11 m 111111 m 11 m 111111111111 n 11111111 ili h II1111 



0 10 20 30 40 60 60 70 80 SO 100 110 120 130 140 150 1601 17 

llliiliiiilliiilmiilimlii ill ii 1111111111111111 m 111 m 111111 m 1111111111111 in 111 n 1111111111111111 m 1111111111 ii 1111111111111111 m 111111111111111111111 ill 11111111 


f y d « p 

0 10 20 30 40 5 0 60 70 • 80 -SO 100 11D 120 130 140 15Q 160 17C 

111| 11 m 1111111111 iJ 1111111 ill ii 11111111111 ill I 11111 n 111 m m 111111 11 j 1 11 m 11.11111 j 11111.111111111111111111111111111111111111T1111111111 ii 11111111 M m 11111111111 ml 


0 ip 20 30 40 50 60 70 00 <0 100 110 120 130 140 150 168 17,( 

1 111 11 1 1 11 1 1111 1111 hi 11 1 1111 1111 1 111 1 111 ill I m 1111111 ill 1111111 hi 1 1 m I m 1 1 1 m 1 1 m 1 1 m 1 1 m I m 1 1 1111 1 1 1 11 1 m 1 1 1111 1 m 1 1 m 1 1 11 ii I m 1 1 m 1 1 m 1 1 11H I ill 11 111111 


I 
















A SHORT TEXT-BOOK 


OF 

INORGANIC CHEMISTRY 


BY 

D R HERMANN KOLBE 

M 

PROFESSOR OF CHEMISTRY IN THE UNIVERSITY OF LEIPZIG 


TRANSLATED AND EDITED BY 

T. &' HUMPIDGE, Ph.D. B.Sc. (Lond.) 

PROFESSOR OF CHEMISTRY AND PHYSICS 
IN THE UNIVERSITY COLLEGE OF WALES, ABERYSTWYTH 


WITH A COLOURED TABLE OF SPECTRA AND 
NUMEROUS WOOD ENGRAVINGS 


NEW YORK 

JOHN WILEY & SONS, 15 ASTOR PLACE 



a ai33 

,K7 




Although the number of short Text books of Inorganic 
Chemistry is large, it is hoped that this little book will supply 
a definite want among teachers and students, corresponding to 
that which the Editor has himself felt. 

The principles which have guided the Author in writing the 
book are fully stated in his Preface, and with these the Editor 
entirely concurs ; but in adapting the book for English students 
certain alterations and additions were necessary, and to these 
the Author has given his full consent. 

The whole book has been carefully revised throughout, and 
the physical constants brought up to date. Considerable 
additions have been made to the descriptions of water, atmo¬ 
spheric air, coal, iron, &c. Short accounts of Gay-Lussac’s 
law, Avogadro’s law, and the manufacture of coal-gas have also 
been introduced into the text. 

In the Appendix, which is entirely new, a brief account has 
been given of the methods used for determining atomic and 


A 2 



iv Editor's Preface. 

molecular weights, of Prout’s law, and of the Periodic law. 
The Editor acknowledges his indebtedness to Lothar Meyer’s 
1 Moderne Chemie ’ in writing this Appendix. P'inally, a series 
of tables has also been added, which it was thought would prove 
useful in the laboratory. Some of these have been taken from 
Landolt and Bornstein’s excellent collection of physical tables. 

The range which the book covers is rather more than that 
required for the Intermediate Science and Preliminary Scien¬ 
tific (M.B.) Examinations of the London University, and the 
needs of students working for these examinations have been 
steadily kept in view, but without following the syllabus in a 
servile manner. 


Aberystwyth : March , 1884 



RS PREFACE. 


This short Text-book has been written to recall to the memory 
of students who have attended a course of lectures on Experi¬ 
mental Chemistry what they have seen and heard, and to clear 
up any points which may not have been properly understood. 

A necessary condition for using a text-book of chemistry 
successfully is attentive and continuous attendance at the 
lectures. Students in arts who have not regularly attended a 
course of lectures may be able to read up afterwards what they 
have missed ; but a lecture which has not been attended by a 
student of chemistry cannot be made up by mere reading— 
neither the notes of the lecture by other students, which ought 
to be extremely few, nor a text-book, can serve as a substitute for 
what has not been heard. The chemist has to learn, not by 
reading nor by hearing alone, but both by hearing and seeing. 
A person who has not seen the phenomena produced by the 
union of oxygen and hydrogen, for example, can have no clear 
conception of them, nor of the chemical change which accom¬ 
panies them. Nothing is more foolish than the opinion, which 


VI 


Author's Preface. 

I have often heard from young medical students, that chemistry 
can be studied from books alone, like other subjects ; that 
facts which are learnt by heart can take the place of general 
principles only partially understood. 

The opinion of chemists as to how many of the immense 
number of empirical facts now known in chemistry should be 
introduced into lectures appears to be very different. I quite 
agree with Fittig when he says : ‘ Chemistry can no longer be 
taught as a descriptive science ; in lectures on chemistry the 
general nature and principles of chemical phenomena should 
take the first place, and should be illustrated by experiments, 
but all the many compounds should not be enumerated and 
described.’ But my opinion differs from that of Fittig when 
he, rather inconsistently, goes on to add : £ But in teaching 
chemistry in this manner it is still requisite to offer to the student 
as easily and completely as possible the material on which 
these general principles are based.’ 

The study of chemistry is similar to that of a language. 
What is learnt in the lectures scarcely goes as far as reading 
and parsing ; at most it only includes the rules by which words 
are built up into sentences. To use the language with success 
—to speak it—continued practice is required. 

The same is also true of chemistry : the science is learnt in 
the laboratory, not in the lecture theatre. The most that can 
be done in the lectures is to prepare the student for successful 
work in the laboratory. 


Author’s Preface. vii 

And although teachers in law, history, and philosophy give 
the best they have in their lectures, experimental chemistry, as 
taught in the lecture theatre, must be elementary. To enume¬ 
rate many chemical facts loads the memory with ballast, and 
tends to learning by rote instead of clear perception and after¬ 
thought. 

The problem of the lecturer on chemistry is therefore to 
give his hearers an idea of chemical processes and the most 
important chemical theories without burdening their memories 
with a large number of mere facts, and thus to prepare them 
to acquire an accurate knowledge of chemistry by their own 
practical work. 

I have adhered to this general principle in writing this short 
Text-book. I have also endeavoured not simply to give a series 
of dry facts, but to blend them together into one continuous 
narrative. 

Of the students attending chemical lectures, those making 
chemistry a special study are nearly always in a minority. A 
complete description of those parts of the science (e.g. the rare 
elements, the ammoniacal compounds of cobalt and the platinum 
bases) which have only interest for these few and not for those 
studying medicine, pharmacy, agriculture, &c., is therefore out 
of place both in lectures and in elementary books ; and all the 
more so as the special study of these subjects is better carried 
on in the laboratory. Such subjects are therefore briefly 
treated in this Text-book, while others of general interest, such 


viii Author’s Preface. 

as water, atmospheric air, carbon, carbonic acid, arsenious acid, 
the detection of arsenic in cases of poisoning, salts, iron, lime, 
&c., are referred to more fully. 

The engravings in the text are limited in number : they 
illustrate only that apparatus which I consider to be especially 
adapted to recall to the student what he has seen in thg 
lectures. 


CONTENTS 


PAGE 


Introductory . . . . i 

Oxygen . 8 

Hydrogen . . . .16 

Combustion.29 

Laws of Chemical Combina¬ 
tion .37 

Law of Multiple Propor¬ 
tions .42 

Law of Volumes . . . 43 

The Atomic Theory and 
Chemical Valency . . 44 

Avogadro’s Law . . .50 

Chemical Affinity . . . 51 

Chemical Nomencla¬ 
ture .54 

Chemical Symbols and For¬ 
mulae .58 

General Remarks on the 
Chemical Elements . . 63 

Physical and Chemical 
Properties of Bodies . . 67 

Crystalline Shape . . 70 

Chemical Properties of 

Bodies.77 

Water.80 

Hydrogen Peroxide . . 90 

Ozone.91 

The Halogens . . .96 


PAGE 

Chlorine.97 

Hydrochloric Acid . . . 104 

Oxides and Oxy-acids of Chlo¬ 
rine .109 

Chloric Acid . . . . no 

Perchloric Acid . . *113 
Chlorine Peroxide . . . 114 

Chlorous Anhydride and Acid 114 


Hypochlorous Anhydride and 


Acid.116 

Bromine .1x8 

Hydrobromic Acid . . . 120 

Bromic Acid . . , . 122 

Hypobromous Acid . . 123 

Iodine.123 

Hydriodic Acid . , . 127 

Iodic Acid . . . .129 

Periodic Acid . . . . 131 

Chlorides of Iodine . *133 

Fluorine.133 

Hydrofluoric Acid . . . 134 


Elements of the Sulphur 
Group . . . . . 137 

Sulphur . . . . . 139 

Sulphuretted Hydrogen . . 145 

Hydrogen Persulphide . . 148 

Oxygen Compounds of Sulphur 148 
Sulphurous Anhydride . . 150 

— Acid ..... 154 
Hydrosulphurous Acid . . I 55 

Sulphuric Acid . . . 155 

— Anhydride .... 162 






X 


Contents. 


Sulphur — continued. page 

Nordhausen Sulphuric Acid . 163 

Sulphuryl Chloride . . . 164 

Chlorsulphonic Acid . . 165 

Nitrosulphonic Acid . . 166 

Disulphuric Acid . . . 166 

Thiosulphuric Acid . . . 167 

Dithionio Acid . . .170 

Polythionic Acids . . . 171 

Compounds of Sulphur and 
Chlorine .... 173 

Selenium .... 174 


Compounds of Selenium . . 175 

Seleniuretted Hydrogen . 177 

Selenious Anhydride and Acid 177 
Selenic Acid . . . . 178 

Sulphides of Selenium . .178 

Chlorides of Selenium . . 179 

Tellurium .... 179 

Elements of the Nitrogen 
Group.180 

Nitrogen.181 

Compounds of Nitrogen and 
Hydrogen .... 182 

Ammonia.183 

Hydroxylamine (Oxyam- 
monia) .... 186 

Amidogen and Imidogen . . 188 

Ammonium . . . .188 

Compounds of Nitrogen and 

Oxygen . . . . 189 

Nitric Acid .... 189 

— Anhydride . . . . 192 

— Oxide .... 192 

Nitrous Anhydride and Nitric 
Peroxide . . . . 195 

— Oxide .... 198 

Atmospheric Air . . . 200 

Compounds of Nitrogen with 
the Halogens . . . 204 

Nitrogen Chloride . . . 204 

— Iodide .... 206 

Phosphorus . . . . 206 

Compounds of Phosphorus 
and Hydrogen . . .211 


Phosphorus —continued. 

PAGE 

Gaseous Phosphoretted Hy¬ 


drogen (Phosphine) . . 

211 

Compounds of Phosphorus 


and Oxygen 

213 

Phosphoric Anhydride . . 

214 

Orthophosphoric Acid . 

216 

Pyrophosphoric Acid . . 

219 

Metaphosphoric Acid . 

220 

Phosphorous Anhydride and 


Acid. 

221 

Hypophosphorous Acid 

222 

Compounds of Phosphorus 


and the Halogens . . . 

223 

Phosphorus Trichloride 

224 

— Pentachloride . . . 

225 

Phosphoric Oxychloride 

226 

Bromides of Phosphorus . 

227 

Iodides of Phosphorus . 

228 

Compounds of Phosphorus 


and Sulphur 

228 

Arsenic. 

229 

/ Arseniuretted Hydrogen 


(Arsine) . . . , 

231 

Compounds of Arsenic and 


Oxygen .... 

233 

Arsenious Anhydride and Acid 

233 

Detection of Arsenic in Cases 


of Suspected Poisoning 

237 

Arsenic Anhydride and Acid . 

244 

Compounds of Arsenic with 


the Halogens . . . 

246 

Arsenious Chloride, Bromide, 


and Iodide . ... 

247 

— Fluoride .... 

248 

Compounds of Arsenic and 


Sulphur . . . . 

248 

Arsenic Disulphide and Tri¬ 


sulphide .... 

248 

— Pentasulphide . . . 

250 


Antimony .... 250 
Antimoniuretted Hydrogen 
(Stibine) . . . . 252 

Compounds of Antimony and 
Oxygen . . . .253 

Antimony Trioxide . . . 253 




Contents. 


xi 


Antimony — continued page 

Antimonic Anhydride (Anti¬ 
mony Pentoxide). . . 254 

— Acid ..... 254 

Metantimonic Acid . . 255 

Antimony Tetroxide . . 256 

Compounds of Antimony and 
the Halogens . . . 256 

Antimony Trichloride . . 256 

— Pentachloride . . . 257 

— Tribromide . . . . 257 

Sulphur Compounds of Anti¬ 


mony . . . . . 258 

Antimony Trisulphide . . 258 

— Pentasulphide . . . 259 

Boron.260 

Boric Anhydride and Acid . 262 

Boron Trichloride. . . 264 

— Trifluoride . . . . 265 

— Sulphide .... 266 

— Nitride . . . . 266 

Silicon . . . ' . . 266 

Compounds of Silicon . . 268 

Silica and Silicic Acid . . 268 

Silicon Hydride . . . 272 

— Chloride .... 273 

— Fluoride . . . . 274 

Fluosilicic Acid . . . 275 

Silico-chloroform . . . 276 

Silico-formic Anhydride,; 

Silicoxalic Acid . . . 277 

Carbon.277 

Compounds of Carbon . . 286 


Carbon— continued . page 

Carbonic Anhydride or Acid . 286 

Carbamic Acid . . . 293 

Carbonic Oxide . . . 294 

Carbon Oxychloride . . 297 

Oxalic Acid . . . . 297 

Oxamic Acid and Oxamide . 299 

Carbon Disulphide . . . 300 

— Oxysulphide . . . 302 

Compounds of Carbon and 
Hydrogen . . . . 303 

Methane or Marsh Gas . . 304 

Ethylene .... 306 

Carbon Tetrachloride . . 309 

Cyanogen . . . .311 

Hydrocyanic Acid . . . 313 

Cyanic Acid .... 316 

Sulphocyanic Acid . . . 317 

Chlorides of Cyanogen . . 318 

Coal Gas.319 

Titanium .... 322 
Titanic Anhydride and Acid . 323 

Titanium Tetrachloride . . 324 

Fluotitanic Acid . . 325 

Molybdenum . . . . 325 

Molybdic Anhydride and 
Acid ..... 326 

Tungsten .327 

Tungstic Anhydride and Acid 328 

Vanadium . . . . 329 

Niobium and Tantalum . . 330 


THE METALS. 


The Metals . . • • 33 1 

Chemical Constitution of 
Salts . . • • 33 2 

Spectrum Analysis . . . 337 

Classification of the Metals . 340 

Metals of the Alkalies . 342 

Potassium .... 342 

Compounds of Potassium . 346 

Potassium Oxide and Per¬ 
oxide . 346 


Pot A ss 1 um —contin tied . 


Potassium Hydrate , 

• • 346 

— Sulphate . 

• 349 

— Acid Sulphate 

• • 349 

— Nitrate 

• 350 

Gunpowder 

• • 35 i 

Potassium Nitrite . 

• 352 

— Carbonate . 

• • 353 

— Acid Carbonate 

• 354 

— Chlorate 

• • 354 

— Perchlorate 

• 354 







Contents. 


xii 


Potassium— continued. 

PAGE 

Potassium Hypochlorite . • 

354 

— Oxalate .... 

355 

— Silicate . . . . 

355 

— Chloride .... 

35 6 

— Bromide . . . . 

356 

— Iodide .... 

356 

— Fluoride . . . . 

358 

— Acid Fluoride . 

358 

— Fluoborate . . . . 

358 

— Fluosilicate 

359 

— Cyanide . . . . 

359 

— Cyanate .... 

361 

— Sulphocyanate . . . 

361 

Sulphur Compounds of Po¬ 


tassium 

362 

Potassium Monosulphide . 

362 

— Sulphydrate 

363 

— Trisulphide and Penta- 


sulphide . 

363 

Liver of Sulphur . 

363 

Detection of Potassium Com¬ 


pounds . 

364 

Sodium. 

365 

Compounds of Sodium . . 

367 

Sodium Oxide and Peroxide . 

3 6 7 

— Hydrate . . . . 

367 

— Sulphate .... 

367 

— Acid Sulphate . . . 

368 

— Sulphite .... 

368 

— Acid Sulphite . . . 

368 

— Thiosulphate . 

368 

— Nitrate . . . . 

369 

— Carbonate 

369 

— Acid Carbonate . . . 

374 

— Phosphates 

374 

— Borate. 

376 

—: Chloride .... 

376 

Detection of Sodium Com¬ 


pounds . 

378 

Lithium . 

378 

Lithium Hydrate . . . 

379 

— Sulphate .... 

379 

— Phosphate . . . . 

379 

— Carbonate 

380 

— Chloride . . . . 

380 


Lithium —con tin ued. 


PAGE 

Detection of Lithium 

Com- 


pounds 

• 

380 

Rubidium and Cesium 

. • 

380 

Ammonium 


382 

Ammonium Sulphate 

. 

382 

— Nitrate 

. 

383 

— Phosphates . 

. 

383 

— Carbonate 

• 

383 

— Chloride 

. 

385 

— Sulphides . 

. 

386 

Oxyammonium Compounds . 

386 

Detection of Ammonium 


Compounds 

* 

387 

Metals of the Alkaline 


Earths 


388 

Calcium . 

. 

388 

Calcium Oxide . 

. 

389 

— Hydrate . >. 


390 

— Sulphate 


394 

— Nitrate 

. 


— Phosphate . 


396 

— Carbonate 

. 

398 

Bleaching Powder . 


400 

Calcium Silicates . 


401 

Glass 


401 

Calcium Oxalate . 


404 

— Chloride 


4°4 

— Fluoride . 


405 

— Sulphides 


405 

Detection of Calcium 

Com- 


pounds . 

. 

406 

Strontium 


406 

Strontium Oxide 

. 

406 

— Sulphate . 

, , 

407 

— Nitrate 


407 

— Carbonate 


407 

— Chloride 

• # 

407 

--Sulphide . 

• t 

407 

Detection of Strontium Com- 


pounds 


408 

Barium 


408 

Barium Oxide 

. , 

408 

— Hydrate 


409 

— Peroxide . 


410 

— Sulphate 

. 

411 





Contents. xiii 


B AR ium— continued. 

PAGE 

Aluminium— continued. 


PAGE 

Barium Nitrate 

4 i i 

Ultramarine 

. 

427 

— Carbonate . . . 

411 

Aluminium Chloride 


428 

— Chloride .... 

411 

— Fluoride 

. 

428 

— Sulphides . . . . 

412 

— Sulphide . 


428 

Detection of Barium Com¬ 


Detection of Aluminium Com- 


pounds ..... 

412 

pounds 


429 

Magnesium .... 

4 i 3 

Bfryllium . 

. 

429 

Magnesium Oxide . . . 

414 

Beryllium Oxide and Hydrate 

430 

— Hydrate .... 

4 i 5 

— Sulphate 


430 

— Sulphate . . . . 

415 

— Chloride . 


430 

— Phosphate? 

416 

— Carbonate 


430 

— Carbonate . . . . 

4 i 7 

Gallium .... 


431 

— Borate .... 

417 

Indium 


432 

— Silicate..... 

417 

Indium Oxide 


432 

— Chloride .... 

4 T 7 

— Sulphate 


432 

Detection of Magnesium 


— Chloride . 


432 

Compounds 

418 

— Sulphide 


432 

Metals of the Earths . . 

419 

Yttrium, Terbium, Erbium, 


Aluminium .... 

419 

and Ytterbium 

• • 

433 

Aluminium Oxide . . . 

421 

Cerium, Lanthanum, 

AND 


— Hydrate .... 

421 

Didymium . 

. . 

433 

— Salts ..... 

422 

Thorium . . . 


434 

— Sulphate .... 

422 

Zirconium . 

. 

434 

Potash-alum . . . . 

423 

Zirconium Oxide and Hydrate 

435 

Aluminium Phosphate . 

425 

— Sulphate 


435 

— Silicate . . . . 

425 

— Chloride . 


435 

Porcelain and Earthenware . 

426 

— Fluoride 

• 

435 

HEAVY 

METALS. 



Iron . 

436 

Compounds of Iron 

AND 


Ferrous Oxide and Hydrate . 

445 

Cyanogen 


453 

— Sulphate . . . . 

445 

Potassium Ferrocyanide 


453 

— Carbonate 

446 

Ammonium Ferrocyanide 

• 

456 

— Phosphate . . . . 

447 

Sodium Ferrocyanide 

. 

456 

Ferric Oxide and Hydrate 

447 

Prussian Blue 

. 

456 

— Sulphate . . . . 

448 

Hydroferrocyanic Acid 

• 

457 

— Nitrate .... 

449 

Potassi um Ferricyanide. 

. 

457 

— Phosphates . . . . 

449 

Turnbull’s Blue 


458 

Triferric Tetroxide 

449 

Hydroferricyanic Acid . 

. 

458 

Ferric Acid . . . . 

450 

Sodium Nitroprusside 

• 

458 

Ferrous Chloride . 

450 

Detection of Iron Compounds 

459 

Ferric Chloride. . . . 

45 i 




Ferrous Iodide 

45 i 

Manganese 

• • 

460 

— Sulphide . . . . 

452 

Manganous Oxide and 

Hy- 


Ferric Disulphide . 

452 

drate . 

. . 

461 







XIV 


Contents. 


A a N G A N E s E —con tin ued. 

PAGE 

Manganous Sulphate 

461 

— Nitrate . . . • 

462 

— Carbonate 

462 

— Chloride . . . • 

462 

— Sulphide .... 

4 6 3 

Manganic Oxide and Hydrate 

463 

— Sulphate .... 

463 

— Chloride . . . • 

464 

Trimanganic Tetroxide . 

464 

Manganese Peroxide . . 

464 

Manganic Acid 

465 

Potassium Manganate 

465 

Permanganic Acid 

466 

Potassium Permanganate . 

466 

Detection of Manganese Com¬ 


pounds . ... 

476 

Chromium .... 

468 

Chromic Oxide. . . . 

469 

— Hydrate .... 

470 

— Sulphate . . . . 

47 i 

— Nitrate .... 

472 

— Phosphate . . . . 

472 

— Chloride .... 

472 

Chromous Chloride . . . 

473 

Chromic Acid 

473 

Chromates . . . . 

474 

Potassium Dichromate . 

475 

— Chromate . . . 

476 

Sodium Chromate . 

477 

Ammonium Chromate and 


Dichromate 

477 

Chlorchromic Acid and Chro- 


myl Chloride 

477 

Detection of Chromium Com¬ 


pounds . 

479 

Uranium - . . . . 

479 

Uranic Nitrate 

480 

— Oxide. 

480 

— Sulphate . . . . 

480 

— Carbonate . . . . 

481 

— Phosphate 

481 

Uranyl Chloride . . . 

481 

— Sulphide . 

482 

Uranous Oxide. . . . 

482 

— Chloride . 

482 

Uranoso-uranic Oxide . . 

482 


Cobalt and Nickel 

Cobalt . 

Cobaltous Oxide and Hydrate 
Cobaltic Oxide and Hydrate . 
Cobaltous Sulphate 
— Nitrate . . • • 

— Nitrite .... 
— Phosphate . . . • 

— Carbonate 

— Oxalate . . . • 

— Chloride .... 
Cobaltic Chloride and Am¬ 
monia .... 
Cobaltous Cyanide. . . . 

Potassium Cobalticyanide 
Cobalt Sulphides . . . 

Smalt . . • . 

Detection of Cobalt Com¬ 
pounds .... 

Nickel. 

Nickelous Oxide and Hydrate 
Nickelic Oxide and Hydrate . 
Nickel Sulphate 
— Nitrate. . . . • 

— Carbonate 

— Chloride . . . . 

— Cyanide .... 
— Sulphide . . . • 

Alloys of Nickel 
Detection of Nickel Com¬ 
pounds .... 

Zinc. 

Zinc Oxide .... 
— Hydrate . . . . 

— Sulphate .... 
— Nitrate . . . . 

— Carbonate 

— Chloride . . . . 

— Sulphide .... 
Detection of Zinc Compounds 

Cadmium . . . • 

Compounds of Cadmium . . 

Detection of Cadmium Com¬ 
pounds. . . . . 

1 Lead . 


PAGE 

483 

484 

485 

485 

486 
486 
486 
486 

486 

487 

487 

487 

488 
488 

488 

489 

489 

490 

490 

491 
491 
491 
491 
491 

491 

492 
492 

492 

493 

494 

495 
495 
495 ' 
495 

495 

496 

496 

497 

497 

498 
498 







Conte?its. 


Lead —contin ued. 

Lead Suboxide . 


PAGE 

501 

— Oxide 

\ 

501 

— Hydrate 

. 

502 

Red Lead 

• 

502 

Lead Peroxide . 

, 

502 

Salts of Lead 

. 

503 

Lead Sulphate . 


503 

— Nitrate . 


504 

— Phosphate . 


504 

— Silicates . 


504 

— Carbonate 


504 

White Lead . 

. # 

504 

Lead Acetate . 

. , 

505 

— Chromate . 


506 

— Chloride 

# , 

S06 

— Iodide 

• v 

S06 

— Sulphide 

• 

507 

Alloys of Lead 

• 

5 °7 

Detection of Lead 
pounds . . 

Com- 

508 

Thallium . 

. 

509 

.Thallous Compounds 

. . 

5 io 

Thallic Compounds . 


5 12 

Detection of Thallium Com- 


pounds . 

• * 

512 

Bismuth . 

t . 

5 i 3 

Bismuth Suboxide . 

, # 

5 i 4 

— Oxide and Hydrate 

• . 

5 i 4 

Bismuthic Acid 


5 i 4 

Bismuth Nitrate . 

• 

5 i$ 

— Chromate 


5 i 5 

— Chloride . 


5 i 5 

— Sulphide 


5 i 5 


Bismuth—< wz tin ued . 

Detection of Bismuth Com¬ 
pounds 

Tin .... 

Stannous Compounds 
— Oxide and Hydrate 
— Chloride 
— Sulphide . 

Stannic Compounds . 

— Oxide and Hydrate 
— Chloride 
— Sulphide . 

Alloys of Tin 
Detection of Tin Compounds 


Copper 

Cupric Compounds 
— Oxide . 

— Hydrate . 

— Sulphate 

— Ammonium Sulphates 
— Nitrate 
— Phosphate 
-— Arsenite 
— Carbonates 
— Acetates 
— Chloride . 

— Bromide 
— Sulphide . 

Cuprous Compounds 
— Oxide and Hydrate . 
— Chloride 
— Iodide 
— Sulphide 
Alloys of Copper . 
Detection of Copper . 


NOBLE METALS . 


Mercury . 

Mercuric Compounds 
— Oxide . 

— Sulphate . 

— Nitrate. 

— Chromate . 

— Chloride 
— Iodide 
— Cyanide 


533 

535 

535 

535 

53 6 
536 
536 
538 
538 


Mercury— continued . 
Mercuric Sulphide. 
Mercurous Compounds 
— Oxide 
— Nitrate 
— 1 Chromate . 

— Chloride 
— Bromide . 

— Iodide . 


XV 


PAGE 

515 

517 
5^7 

518 

518 

519 

519 

520 
520 
521 

C22 

523 

524 

524 

525 

525 

526 
527 
527 
527 

527 

528 
528 

528 
528 

529 
529 

529 

530 
530 

530 

531 


538 

539 
539 

539 

540 

540 

54 1 
54 i 













XVI 


Contents. 


Mercury — continued. page 

Mercurous Sulphide . . 54 1 

Amalgams . . • 54 1 

Detection of Mercury Com¬ 
pounds . . . • S 4 2 

Silve t . 543 

Silver Oxide . • • 545 


— Peroxide .... 54 ^ 

— Sulphate . . • 54 ^ 

— Nitrate .... 546 

— Phosphate . . • • 547 

— Carbonate . . • 547 

— Chromate . . * • 54 8 

— Chloride. Bromide, and 

Iodide . . . .548 

— Fluoride . . . • 55 ° 

— Cyanide .... 55 ° 


Silver— continued. page 

Silver Sulphide . . • 55 ° 

Alloys of Silver . . • • 55 1 

Detection of Silver Com¬ 


pounds . . • • 55 1 

Gold. 55 2 

Auric Chloride . . • 553 

— Oxide and Hydrate . . 554 

Potassium Aurate . . • 554 

Auric Cyanide . . • 554 

— Sulphide . . • • 555 

Aurous Chloride . • • 555 

— Oxide . . ... 555 

— Cyanide .... 555 

Purple of Cassius . . . 55 6 

Alloys of Gold . . • 55 ^ 


PLATINUM METALS. 


Platinum .... 55 8 
Platinic Chloride . . • 5 61 

Platinous Chloride . • 5 ^ 2 

Platinic Oxide and Hydrate . 562 

Platinous Oxide and Hydrate 563 
Platinic Sulphide . . . 563 

Platinous Sulphide . . 5^3 

— Cyanide . . • • 5^3 

Ammoniacal Compounds of 

Platinum . . . • 5^3 

Palladium .... 566 

Palladium Hydride . . . 566 


Pall adi u m —contin 7 ied . 

Palladous Nitrate . . . 567 

Palladic Chloride . . . 568 
Palladous Chloride . . 568 

— Iodide ... . • 568 

Iridium.5 68 

Compounds of Iridium . . 569 

Rhodium. 57 ° 

Osmium. 57 ° 

Osmic Acid .... 570 

Ruthenium . 57 1 


APPENDIX. 

Atomic and Molecular Weights.573 

Relations between the Atomic Weights.582 

Tables.5 8 6 

INDEX . . . . ,. 593 


TABLE OF SPECTRA . 


Frontispiece ■ 












INTRODUCTORY. 

The science of Chemistry is closely related to that of Physics, and 
it.is the aim of both sciences to solve nearly the same problems. 
The province of both is to investigate the processes of nature per¬ 
ceptible to the senses, to discover the course they take, their con¬ 
nection with one another, and their causes, and to investigate the 
laws according to which the forces of nature act, whatever may be 
the objects or classes of bodies upon which their action is exerted. 

These two sciences—Chemistry and Physics—are the foundation 
of all other natural sciences—of mineralogy and geology, of zoology 
and botany, as well as of astronomy, and especially of the whole of 
medical science, which endeavours to fathom the processes pro¬ 
duced in a certain class of bodies by chemical and physical forces. 

If we consider what goes on around us in nature, we perceive 
the change from day to night, we notice the moon and apparently 
the sun revolving round the earth, as well as similar movements in 
other heavenly bodies ; we feel the movements of the air, perceive 
the change of temperature ; we notice water freezing, and admire 
the rainbow, the thunderstorm, formation of clouds, &c. All these 
phenomena belong to the processes of physics ; but where shall we, 
then, look for chemical processes ? 

Simple chemical processes are not nearly so common in nature 
as simple physical processes, and this is certainly the reason why 
physical phenomena and physical laws were investigated many 

B 


X 





2 Text-Book of Inorganic Chemistry. 

centuries before any idea of a natural law in chemistry was con 
ceived. 

We notice plants growing and animals breathing, both which 
processes go on by the inter-action of chemical and physical forces. 
But how plants assimilate the constituents of atmospheric air, and 
work them up chemically into their substance, still remains a 
mystery to the chemist and physicist at the present day. 

When a tree is struck by lightning and burns, this is certainly 
a chemical process, and one much simpler than the growth of a 
plant. But although man has known from time immemorial how 
to produce fire, and has had the chemical process of burning wood 
daily before his eyes, a correct explanation of a process apparently 
so simple was only obtained at the end of the preceding century. 
For this reason chemistry may well be called the younger sister of 
physics. 

The intelligent observer also perceives numerous other chemi¬ 
cal processes going on in nature ; he notices the weathering of 
rocks, the putrefaction of animal and vegetable matter, the deposi¬ 
tion of limestone in caverns, the rusting of iron, the souring of 
milk, wine, &c. ; but these and similar changes usually go on so 
slowly that their progress is difficult to follow. 

Innumerable chemical processes were introduced by man 
thousands of years ago. It was found that lime which had lain in 
the fire—quick-lime—became hot when brought into contact with 
water, that it possessed generally other properties than the lime 
which had not been burnt; metals had been extracted from their 
ores by heating with charcoal; white arsenic had been obtained 
by roasting other ores ; sulphur had been burnt in the air and the 
suffocating acid gas so produced had been noticed. But these and 
a hundred other similar changes could only be correctly understood 
and explained after the common cause of them all—fire—and the 
process of combustion had been made clear. 

For this reason, the discovery of oxygen—of that substance 
which is as necessary for combustion as it is for respiration—by 
Priestley and by Scheele, and Lavoisier’s first correct explanation 
of combustion, are the pillars which carry the noble edifice of 
modern chemistry, based on the foundation of numerous earlier 
observations. 

In what manner, then, do chemical and physical processes differ 
from one another—which phenomena belong to chemistry and 
which to physics ? The answer may be easily given in a few words, 


Introductory. 3 

but will not be so easily understood by those to whom chemical 
phenomena are strange. 

Chemical processes are those by which bodies undergo a material 
change. In physical processes, pure and simple, the substances 
in which we notice a change remain materially unchanged. But 
what is to be understood by a material change ? This will best 
be explained by a few examples. 

A stick of sulphur when rubbed becomes electric and attracts 
small pieces of paper or other light bodies, and then repels them. 
The same sulphur, if heated in a test tube, melts to form a pale 
yellow, clear liquid, and changes at a higher temperature to a 
reddish-brown gas, which externally has not the remotest similarity 
with the solid yellow sulphur. But has the sulphur by the rubbing 
or by the heating undergone a material change, have chemical pro¬ 
cesses taken place in both cases ? 

The answer to this question is a negative. The substance of the 
sulphur is the same after the rubbing as before, and sulphur gas, 
as well as liquid sulphur, have the same relation to solid sulphur 
as water gas or water vapour, and liquid water to solid water or ice. 
We have here to do with a change in the state of aggregation pro¬ 
duced by heat, which only lasts as long as the cause producing it. 
Just as water gas by cooling again becomes first liquid water and 
then solid water, so the sulphur gas changes into liquid sulphur 
and then into solid sulphur, with exactly the same properties as it 
possessed at first, as soon as the source of heat is removed. 

But something quite different happens when we heat sulphur 
in an open vessel, so strongly that it catches fire and burns with a 
blue flame. In this case, too, it changes entirely into a gas, the 
colourless gas with the well-known odour of burning sulphur. If 
we cool the gas, for example, by leading it through a vessel sur¬ 
rounded with cold water, we should in vain expect to obtain liquid 
or solid sulphur from it. By this process the sulphur has suf¬ 
fered a material change— i.e. a change which continues after the 
cause has been removed. The sulphur has united chemically 
with a constituent of the air, and a new substance has been 
produced. 

If we intimately mix together yellow sulphur and metallic iron, 
both in the state of a fine powder, we obtain a grey substance, 
which externally has no similarity with either of the bodies from 
which it was produced. It might be thought that this grey powder 
no longer consists of yellow sulphur and black iron, and that the 


4 


Text-Book of Inorganic Chemistry. 

two substances had undergone a material change by the act of 
mixing them. 

We can, however, readily convince ourselves that this mixture 
really contains unchanged sulphur and iron. The lighter sulphur 
may be easily separated by the mechanical process of washing, and 
the iron may be removed by a magnet—which would not be pos¬ 
sible with the product of real chemical action. 

Let us prepare a mixture of four parts of sulphur and seven parts 
of iron, both finely powdered, and heat the mixture in a test-tube 
in the gas lamp. In a short time that portion 6f the mass which 
is most strongly heated commences to glow'; the glowing then 
spreads through the entire mixture, and finally ceases of itself. It 
seems remarkable that at this high temperature none of the volatile 
sulphur is given off. 

Sulphur and iron, which wdien only mixed remain chemically 
unchanged, attract one another chemically when heated to a cer¬ 
tain high temperature, and then both of them suffer a material 
change. The dark-grey, solid, hard product may be easily pulver¬ 
ized, but this powder has now quite different properties to the mere 
mixture of sulphur and iron. We might try in vain to separate 
the sulphur by washing or by any other mechanical operation, or 
to extract the iron by the magnet. Sulphur and iron experience a 
material change when heated together, and in such a manner that 
it might be asked whether the new' product really contains sulphur 
and iron, or whether both have not been lost by the operation. 

This question cannot be decided a priori, nor by mere philoso¬ 
phising, but only by experiment, and it has been so decided. 

If the sulphur and the iron had been lost on their union, if 
they were no longer contained in the compound, it would not be 
possible in any imaginable manner to again separate them with 
their original properties. But chemistry teaches that, in this and 
other cases, compounds may be again decomposed into their con¬ 
stituents, and that the same quantities of these constituents may be 
again obtained. The process of building up a chemical compound 
from its constituents, as here illustrated in the case of sulphur and 
iron, is called synthesis : the reverse of this is analysis. 

The union, or, as we may say, the coalescence—of the consti¬ 
tuents of a chemical compound is so close that by no means can 
either of them (eg. the yellow sulphur in the above compound) be 
perceived in the dust of a compound, not even wfith the most 
powerful microscope. It is thus clear why it is impossible to split 


Introductory. 5 

up a chemical compound by pure mechanical means. Substances 
united by chemical force can be only separated by some force acting 
chemically . 

Those substances which we cannot further decompose by 
chemical analysis we call simple bodies, or elements , without wish¬ 
ing to assert that they are really undecomposable. They are 
elementary substances for us only as long as we have not succeeded 
in decomposing them. 

The inquiry after the fundamental substances of inorganic and 
organic nature is as old as the investigation of nature itself; but 
what we understand to-day by the word chemical element is some¬ 
thing quite different from what was formerly so called. The state¬ 
ment of the early philosophers, who neglected experiment and 
exact investigation, that Fire , Air , Water , and Earth were elements, 
rather meant that all natural bodies have their origin in these four 
things, than that they consist of them. And when this question 
was more critically examined and experimentally investigated, 
these four elements of the ancients were abandoned. 

It is now known that what is called fire is not a substance but a 
phenomenon; that water consists of two gaseous bodies, as yet 
incapable of further decomposition, oxygen and hydrogen ; that 
air is a mechanical mixture of two gaseous substances, oxygen and 
nitrogen ; and that earth is no single individual substance, but 
an agglomerate of thousands of very different chemical elements 
and compounds. 

Our present chemical elements are the result of experiment. 
Their number is somewhat large, and, according to experience, 
will probably increase. We already know more than sixty. 

When it is considered how easy it is for us to decompose water 
into its constituents, and to reproduce it from them, the question 
arises, how was it possible that up to a hundred years ago water 
was thought to be an elementary substance, how was it possible 
that one of the many methods which we now know for its decom¬ 
position, had not been discovered much earlier ? 

The principal agent, which was previously almost exclusively 
employed in order to produce chemical decomposition, was a high 
temperature, or fire. Crucibles, heating furnaces, distilling appa¬ 
ratus were most indispensable in furnishing the laboratory of an 
alchemist. If water had possessed the property to be decomposed 
at a high temperature, its compound nature would have been re¬ 
cognized long previously. That it remains unchanged at a red- 


6 Text-Book of Inorganic Chemistry . 

heat could only help to confirm the belief that it was undecompos- 
able. 

The methods which we now use to decompose water were 
partially unknown to the chemist of earlier centuries, and were 
partially wrongly interpreted. 

To the former, belong the electrolytic force of the galvanic 
current, and the action of those metals, only discovered later, which 
decompose water at the ordinary temperature : potassium, sodium, 
&c. To those processes by which water is decomposed, and which 
were long known, but were previously falsely explained, belongs the 
reaction which occurs when iron is brought into contact with water 
and a strong acid. Everyone who sees this experiment for the 
first time believes, as was formerly done, that the gas evolved from 
the iron is extracted from this metal, whereas it really arises from 
the decomposition of the water. 

Let us consider this and other processes by which water is de¬ 
composed more exactly. 

If we dip two platinum plates into water which is made slightly 
acid with sulphuric acid (pure water is a bad conductor of the cur¬ 
rent), and then 'connect the plates with the poles of a galvanic 
battery of four or six Bunsen’s cells, we notice at the moment when 
the circuit is completed that gases are evolved from both platinum 
plates, and that this evolution continues as long as the galvanic 
current remains unbroken. 

If two small tubes of equal size, filled with water, are placed 
over the platinum plates, so that the gas evolved from one plate 
is collected in the one tube, and that evolved from the other plate 
in the other (fig. i), we notice the following results 

i. The gas collecting over the positive electrode 1 regularly fills 
a smaller space than that collecting over the negative electrode ; 
exact measurements have shown that the volume of the latter is 
just twice as much as that of the former. 

ii. If a lighted body is brought near the gas from the negative 
pole, when it is allowed to flow out by opening the stopcock, it 
catches fire and burns with a feebly luminous flame. The gas 
from the positive pole does not possess this property ; it does not 
burn itself, but it supports the combustion of other bodies in such 

1 The two pieces of platinum which bring the current to the water are 
called poles or electrodes . The one which is connected with the zinc of the 
battery is the negative pole or electrode, the other which is attached to the 
copper, carbon, or platinum of the battery is positive pole or electrode.—E d. 


Introductory. y 

a manner that a glowing chip of wood brought into contact with it 
at once catches fire and burns with a bright light. 

We call the latter gas oxygen , the former hydrogen. These are 
the elementary constituents of water, on the re-union of which 
water again results. 



Fig. i. 


We will now consider each of these substances separately. 
Oxygen may well be taken first, as it belongs to the most widely 
distributed elements, and since our present scientific chemistry 
dates from its discovery, with the correct explanation of the pro¬ 
cesses of combustion which immediately followed. 










8 


Text-Book of Inorganic Chemistry. 


OXYGEN. 

Chemical Symbol : O .—Atomic Weight'. 16. 

Oxygen belongs to those elements which are most widely dis¬ 
tributed in nature. The atmosphere contains 23 per cent, of 
oxygen by weight, water 88*9 per cent., and it is further one of the 
chief constituents of the innumerable chemical compounds which 
make up the solid crust of the earth. Oxygen is as widely dis¬ 
tributed in organic as in inorganic nature ; and it is present.in many 
of the numerous products of animal and vegetable life. Sugar, 
cellulose, starch, and many other similar bodies, contain more than 
50 per cent, of oxygen. 

At the first glance it seems almost incredible that a substance 
so very widely distributed in nature as oxygen, and even occurring 
in the free state in atmospheric air, should not have been dis¬ 
covered earlier than the end of the preceding century. 

The honour of its discovery belongs to two chemists, Priestley 
and Scheele, who, quite independently of one another, observed 
and described it in the year 1774. 

Oxygen is a gaseous substance which can only be condensed 
to a liquid when exposed to immense pressure and extreme cold. 
It was long thought to be only capable of existing in the gaseous 
state, and was one of the so-called ‘ permanent gases.’ Recent 
experiments have shown that, by improved methods of producing 
a high pressure combined with intense cold, oxygen may be 
liquefied. 

The gas is colourless and transparent, without taste or smell, 
and may be respired like common air. It is in fact the free 
oxygen in the air which supports respiration. 

Like common air, it is only slightly soluble in water : 100 
volumes of water free from air at 4 0 dissolve only 37 volumes of 
oxygen, or one kilogramme of water dissolves 0*053 gramme of 
oxygen. 

Oxygen is a little heavier than common air; its specific gravity, 
compared with air as unity, is 1*1056. Since one litre of air at o°, 
and under a barometric pressure of 760 mm., weighs 1*293 gramme, 
one litre of oxygen under similar circumstances weighs 1*429 
gramme. 


Oxygen . 9 

Notwithstanding its immense distribution in nature and its 
innumerable compounds, only a few substances are adapted for 
the preparation of oxygen. 

We may obtain it from water by electrolysis in the manner 
just described. But although this method gives very pure oxygen, 
it is not suited for the preparation of large quantities. Among 
the numerous minerals containing oxygen very few can sene 
directly for its preparation. The foremost of these few is pyro- 
lusite, an oxygen compound of a metal—manganese—closely 
resembling iron. This compound is called by chemists manganese 
peroxide (black oxide of manganese). It is a greyish-black ore 
which occurs somewhat largely in nature, and possesses the 
property when heated of parting with a portion of its oxygen, 
forming another compound of the same metal containing less 
oxygen. 

In order to prepare oxygen from it, a long iron tube, closed at 
one end, about the size of an ordinary gun-barrel, is filled with 
small pieces of the broken mineral. The open end is then closed 
with a cork, pierced to receive a glass tube, and heated in a gas 
or charcoal furnace. As soon as it commences to become red- 
hot, gas is freely evolved from the glass tube, which is. allowed to 
dip under water. The gas which is first given off, and which is 
contaminated with the atmospheric air present in the tube, is 
allowed to escape ; the oxygen may then be collected in glass 
cylinders or in a gasometer as required. 

This method gives large quantities of oxygen easily and quickly, 
but not of perfect purity. A compound of manganese, containing 
less oxygen (trimanganic tetroxide), remains behind in the tube, 
and gives off no further oxygen, even when very strongly heated. 

A larger quantity of oxygen may be obtained by heating the 
mineral with sulphuric acid, but for other reasons this method is 
not to be recommended. 

In order to prepare oxygen from the atmospheric air, the nitro¬ 
gen, with which it is mechanically mixed, must be removed in the 
shape of some non-volatile chemical compound. This, however, 
cannot be done, since the force which produces chemical union, 
and which we call chemical affinity, is very powerful in the case of 
oxygen, and the reverse in that of nitrogen ; in consequence of 
this, nearly all attempts to fix the nitrogen of the air in a chemical 
compound result in binding the oxygen, while the nitrogen remains 
free. 


io Text-Book of Inorganic Chemistry. 

We can, however, accomplish our purpose in a roundabout way. 
We can cause the oxygen of the air to unite with some other body, 
producing a compound, which, like the pyrolusite, again gives up 
its oxygen under favourable circumstances. This may be done in 
the following manner. 

The metal mercury possesses the property of combining chemi¬ 
cally with the oxygen of the air, when heated up to a certain 
temperature (320°), which must not be exceeded to any great 
extent. By this process it loses its metallic lustre, and becomes 
changed into a red powder, called red oxide of mercury (mercuric 
oxide). This product has the remarkable property of decomposing 
into its constituents—mercury and oxygen—when heated to a 



Fig. 2. 


temperature a little higher than that necessary for its production. 
In this way, which was first used by Priestley to prepare the gas, 
oxygen may be indirectly obtained from the air, in a state of great 
purity. 

The accompanying figure (fig. 2), shows a simple apparatus for 
the preparation of oxygen from this mercuric oxide, which may be 
also used to determine the quantity of mercury by weight con¬ 
tained in the weighed quantity of mercuric oxide employed. The 
mercuric oxide is contained in the closed end of the tube of diffi- 
cultly-fusible glass, as shown in the figure. Its weight is deter¬ 
mined by weighing the bent tube when empty, and then again 





when the substance has been introduced ; the difference in weight 
gives the quantity of mercuric oxide employed. A cork with a bent 
glass tube is then fitted air-tight in the open end, its other end 
dipping under a cylinder filled with water. If the mercuric oxide 
is now heated by a gas flame, and the heating continued until it has 
all vanished, the oxygen which was previously present in it passes 
over into the glass cylinder standing over water, and the less 
volatile mercury is deposited in the cooler portions, a , of the tube. 



Fig. 3- 


If the tube containing the metallic mercury is again weighed, 
when perfectly cold, and from this the weight of the empty tube 
subtracted, the difference will give the quantity of the mercury 
produced. At the same time the volume of the oxygen may be 
measured and its weight determined from its volume and specific 
gravity. 

If the experiment is carried out with care and with the neces- 














12 


Text-Book of Inorganic Chemistry. 

sary precautions, it will be found that mercuric oxide always yields 
92*6 per cent, of metallic mercury, and 7-4 per cent, of oxygen, or, 
reckoning the oxygen by volume, that 100 grammes of mercuric 
oxide give 92-6 grammes of mercury, and 5,170 c.c., or 5-17 litres 
of oxygen. 

For the preparation of considerable quantities of pure oxygen, 
the best material is an artificial chemical compound very rich in 
this element—potassium chlorate. This salt, which is soluble in 
hot water, and crystallizes out on cooling in small plates, with a 
mother-of-pearl lustre, consists of the metal potassium and the 
gases chlorine and oxygen. It possesses the property of melting 
when heated, and of giving off all its oxygen with apparent boiling 
and frothing, the end product, a compound of its two other consti¬ 
tuents—viz. potassium and chlorine—remaining behind. In this 
manner 100 grammes of potassium chlorate give 39 grammes of 
oxygen, or more than five times as much as that obtained by heat¬ 
ing an equal weight of mercuric oxide. 

This operation is best carried out in a retort of hard glass, as 
shown in fig. 3. The neck of the retort is connected by a moveable 
india-rubber tube with a glass tube which leads the gas into a 
gasometer, and expels the water as it enters. 

In order to prevent the frothing over of the heated mass in the 
retort, the potassium chlorate may be mixed with an equal weight 
of dry pyrolusite (or oxide of copper) before putting it into the 
retort. Such a mixture does not melt when heated, and as the 
solid body mixed with it distributes the heat equally to the entire 
mass, the evolution of gas takes place without frothing, and much 
more quickly. 1 

If large quantities of oxygen are to be prepared, and a pound 
or more of the mixture of manganese peroxide and potassium 
chlorate heated at once, a cast-iron retort may be used, with a 
broad, flat edge (fig. 4). On this edge fits a cover with an enlarge¬ 
ment at the top, and a tube for the escape of the gas. This cover 
is screwed fast, a lute of moist clay free from sand being used to 
make it air-tight. A large gas-burner placed underneath is suffi¬ 
cient to effect the decomposition of the potassium chlorate. 

The chemical properties of a substance are shown by its 
behaviour to other bodies. Let us see how oxygen acts towards 

1 The manganese peroxide remains unchanged during the reaction ; why it 
is that some substances cause the evolution of the gas at a lower temperature 
and others do not, has not yet been satisfactorily explained.—E d. 


Oxygen. 13 

some of the better known chemical substances, and what pheno¬ 
mena may be observed at the same time. 

, Daily experience teaches us that a glowing chip of wood is soon 
extinguished in the air. But if it is dipped into a jar of oxygen it 
not only continues to glow, but at once bursts into flame. A piece 
of charcoal feebly glowing soon goes out in the air, but continues 
to burn with much greater heat and light when introduced into 
oxygen. Sulphur when heated in the air burns with a pale blue 
flame. But if a small piece is placed in an iron spoon, ignited, and 
then dipped into a jar containing oxygen, not only does the flame 
increase in size and brilliancy, but the sulphur also burns away 
more rapidly. We may readily convince ourselves of this by 



Fig. 4. 


burning two pieces of sulphur of the- same size, one in air and one 
in oxygen. Phosphorus, which catches fire so easily, and which 
even in common air burns brilliantly, if heated in an iron spoon 
and introduced into a large jar of oxygen, produces a light so intense 
that the eye cannot bear it, and so high a temperature that the glass 
vessel is often broken. 

Iron, which melts and burns difficultly in the air, catches fire 
when previously heated and plunged into oxygen, burning with 
great brilliancy and a shower of sparks to form a new substance—a 
chemical compound of iron and oxygen. To perform this pretty 
experiment on a small scale it is best to employ a thin steel watch- 
spring. The spring is softened and bent into a spiral form, attached 













14 Text-Book of Inorganic Chemistry. 

at one end to a cork, and to the other end is fastened a small piece 
of tinder. The tinder is then set on fire and the whole plunged 
into a jar of oxygen. The tinder first burns with a bright flame, 
and then imparts its temperature of combustion to the iron, which 

is high enough to make the metal 
burn brightly with a shower of sparks 
(fig. 5). This phenomenon lasts as 
long as any unburnt oxygen remains, 
or until all the iron is consumed. The 
drops of the molten compound of iron 
and oxygen which fall down are so 
hot that they fuse into the bottom of 
the glass vessel, and often cause its 
destruction. 

In the same manner as these sub¬ 
stances, all bodies which burn in 
atmospheric air burn also in oxygen, 
and always with a much greater evo¬ 
lution of light and heat than we are accustomed to see in common 
air. 

In all cases the products of the combustion are chemical com¬ 
pounds of the burnt body with oxygen. With charcoal and sulphur 
these products are not so perceptible as with iron, as they are 
colourless, transparent gases. But their presence may be easily 
made manifest to the eye. The gaseous product of the combustion 
of charcoal, carbonic acid, possesses the property of making a clear 
transparent aqueous solution of lime—lime-water—turbid, by the 
production of a new chemical compound of carbonic acid and lime, 
which is insoluble in water. If, therefore, we pour clear lime-water 
into the jar in which the charc.oal has burnt, and bring it into close 
contact with the carbonic acid by shaking, the liquid becomes 
milky and afterwards deposits a white precipitate. 

The gaseous product (sulphurous acid) of the combustion of 
sulphur may be recognized by its piercing acid odour, as well as 
by the reddening of the blue aqueous solution of litmus when this is 
poured into the jar in which the sulphur has burnt. 

The product of. the combustion of phosphorus—phosphoric 
acid—is really a solid body—a snow-white powder ; it dissolves, 
however, so easily in water that if the vessel in which the phos¬ 
phorus was burnt contained a little moisture the compound at once 
disappears. The presence of this body, or generally that the com- 



Fig. 5- 






Oxygen. 15 

bustion of phosphorus produces a substance with strongly acid 
properties, may also be shown by the strong red colour which is at 
once imparted to blue litmus solution brought into contact with it. 

It will be noticed that the chemical compounds produced by the 
combustion of sulphur and phosphorus are acid bodies ; the same is 
also true for that body produced by the combustion of charcoal. 
At one time it was thought that all acids contained oxygen, whence 
its name, from d£vr,‘ acid,’ and ycwaw, ‘ I produce. 5 We now know 
that this is incorrect. 

The apparatus which served originally for measuring volumes 
of gases, and called, therefore, gasometers, are now principally em¬ 
ployed for the collection and 
preservation of gases. We 
have referred above to these 
instruments, which are pro¬ 
vided with simple arrange¬ 
ments for allowing the gas 
to stream out under pres¬ 
sure of water as required, 
and so make it possible for 
the chemist to use the gas 
at any time. 

The construction of the 
gasometer is simple, and 
may be easily understood 
from the accompanying 
figure (fig. 6), and from 
fig. 3. It consists of two 
vessels of copper (or zinc), 
of which the lower one, B, 
is closed in all directions, 
while the upper and smaller 
one, A, is open at the top. 

Both communicate with two 
tubes, a and 6 , provided Fig. 6. 

with stop-cocks, which, to¬ 
gether with short thick rods c, c, serve to support the upper vessel. 
The tube, b, does not go further than just through the base of A 
and the top of B ; the tube a, on the other hand, extends nearly to 
the bottom of the vessel b, and is open at the lower end. At the 












































































16 Text-Book of Inorganic Chemistry. 

side is another tube with the stop-cock e , which allows the gas 
contained in B to flow out when required. 

The vessel B is first filled with water by pouring it into A, and 
opening the cocks, a , b , and e. The water flows in by the tube a, 
reaching nearly to the bottom of B, and displaces the air, which 
escapes by the tubes b and e, until the entire vessel is full. After 
all the stop-cocks have again been closed, the lower, wider tube d 
maybe opened without water flowing out. The end of the tube, 
from the apparatus evolving the gas, is then inserted at d , and as the 
gas enters the vessel B it displaces the water which flows out by 
the side of the tube from d. 

The extent to which the gasometer is filled with gas is shown 
by the height of the column of water in the glass tube g , f which 
communicates with the interior of the vessel B, above and below, 
and in which the water is always as high as in the interior of the 
gasometer. 

When the gasometer is filled as far as required, the tube is 
withdrawn, the screw at d replaced, and the stop-cock a opened. 
The gas enclosed in B is thus subjected to a pressure equal to the 
column of water from the surface of the water in B to that in A. 

If the stop-cocks and joints of the gasometer are air-tight, the 
gas may be preserved in it for a long time. When a regular stream 
of gas is required, the tube e is connected with the apparatus by a 
piece of india-rubber tubing, and the stop-cock opened. If the gas 
is to be collected in a cylinder over water, the cylinder is filled with 
water, closed with a glass plate, and inverted in the water of the 
vessel A. The glass plate is then removed, and the stop-cock b 
opened, when the gas rises, the cock a being open, under the pres¬ 
sure of the difference between the column of water in and over the 
tube a and of that over the opening of b in the cylinder. 


HYDROGEN. 

Chemical Symbol : H .—Atomic Weight : i. 

Unlike oxygen, hydrogen does not occur in the free state in the 
earth’s atmosphere, but recent physical investigations have shown 
that it is contained in the atmosphere of some of the heavenly 



Hydrogen. 17 

bodies, particularly in that of the sun. As a constituent of water, 
hydrogen is very widely distributed and in very large quantities ; 
it is also contained in nearly all organic compounds, either as water 
or in some other form of combination. Hydrogen owes its name 
to the fact that it is an essential constituent of water (ufiop, ‘ water,’ 
and yevvao), ‘ I produce ’). 

Although hydrogen was prepared as early as the sixteenth cen¬ 
tury by Paracelsus, it was Cavendish who in 1781 first recognized 
it as an elementary body, and described its properties. Cavendish 
may therefore be rightly considered as the real discoverer of this 
element. 

Hydrogen is a colourless gas without smell or taste, which can¬ 
not be respired, and which is even less soluble in water than oxygen. 
It was formerly included under the ‘permanent gases,’ but has now 
been condensed to a liquid \ the pressure and cold required being 
even greater than in the case of oxygen. 

One of the most striking physical properties of hydrogen is its 
low specific gravity. Oxygen, as we have seen, is a little heaviei 
than atmospheric air; hydrogen is so very much lighter that one 
volume of air weighs more than fourteen equal volumes of hydro¬ 
gen. Its specific gravity is exactly 0*0692. One litre of hydrogen 
at o° and under an air pressure of 760 mm. weighs 0*0895 gramme ; 
while under similar conditions one litre of air weighs 1*293 gramme, 
and one litre of oxygen i* 4 2 9 gramme. Air is therefore 1444 
times as heavy as hydrogen, and oxygen 16 times as heavy. 
This property of hydrogen makes it better adapted than any 
other gas for filling balloons. 

Water is principally employed for the preparation of hydrogen. 
The gas may be obtained, as we have previously seen (p. 6), by 
the decomposition of water into its two constituents by the electiic 
current, and by collecting the gases in separate glass tubes, pre¬ 
viously filled with water. Or, we may employ some body acting 
chemically, which, having a stronger affinity for oxygen than 
hydrogen has, unites with the former to produce a non-volatile 
chemical compound, and sets the latter free in the gaseous 
state. 

To those elements which unite a very strong affinity for oxygen 
with a very weak attraction for hydrogen belong the metals potas¬ 
sium and sodium. If a small piece of sodium is thrown upon 
water, it swims on the surface with a hissing noise, evidently pro¬ 
ducing an evolution of gas. The piece gradually becomes smaller 
1 c 


18 Text-Book of Inorganic Chemistry. 

until it completely vanishes. By this reaction the metal sets the 
hydrogen free from the water and unites with the oxygen ; the 
oxygen compound of sodium so produced is then dissolved in the 
remaining water. 

In order to collect the hydrogen evolved, and to recognize it 
as such, a small glass tube, closed at one end, may be filled with 
mercury and inverted in a trough containing the same substance. 
A little water is then introduced into the tube, and a small globule 
of sodium passed up. As soon as the sodium reaches the water 
in the tube a lively evolution of gas commences, the globule be¬ 
comes smaller, and in a few seconds disappears. The gas which 
collects in the upper portion of the tube depresses the column of 
mercury and takes its place. If the tube is now closed with 
the thumb, re-inverted, and a burning chip brought near to the 
open end, the gas, which rapidly streams out on account of its 
lightness, catches fire and burns with a pale, scarcely luminous 
flame. 

This method of obtaining hydrogen is very simple and instruc¬ 
tive, but is not well adapted for the preparation of large quan¬ 
tities of the gas. We are acquainted with other metals that have 
a strong affinity for oxygen, and which can also decompose water, 
which, however, do not, like sodium, possess this property at the 
ordinary temperature, but only acquire it at a higher temperature, 
near a red-heat. To these belongs iron. 

. If an iron tube— e.g. an ordinary gas-tube—is filled with iron wire 
wound together, and then heated to redness in a furnace so that 
the ends project out for some distance, and if steam from boiling 
water in a retort is passed in at one end of the tube, the water is 
decomposed, its oxygen uniting with the iron to form a solid com¬ 
pound which incrusts the metal, while its hydrogen is set free from 
the other end of the tube. If a glass tube dipping under water is 
previously connected with this end of the iron tube, large quanti¬ 
ties of hydrogen can be collected in inverted glass cylinders placed 
to receive.it. 

Iron acquires the property of decomposing water at the ordinary 
temperature, when a strong acid— eg. sulphuric acid—is added to the 
water. Since, however, iron always contains carbon, and often 
other substances mixed with it, the hydrogen so obtained from this 
metal is always rendered impure from the admixture of other sub¬ 
stances, particularly of gaseous hydro-carbons. It is therefore 
better to employ metallic zinc, which is easier to obtain in a fairly 


Hydrogen . I g 

pure state, and which, in this respect, behaves in just the same 
manner as iron. 

As a vessel for generating the gas, a WoulfPs bottle with two 
necks may be employed, which contains the zinc in the granulated 
state (fig. 7). In the one neck is a cork, fitting air-tight, which is 
pierced to receive a funnel-tube, passing to the bottom of the flask, 
and serving to pour in the acid. The other neck also contains a 
cork, carrying a glass tube just passing through it, by which the 
gas escapes and may be collected. For convenience in movement 
this glass tube is best divided into two parts connected together by 
an india-rubber tube. If now sulphuric acid strongly diluted with 
water is poured in by the funnel-tube, and so brought into contact 



Fig. 7 . 

\ 

with the zinc, a copious evolution of hydrogen is at once produced 
with the development of heat. In the flask the salt called white 
vitriol (zinc sulphate) remains behind dissolved in the water. 

Before commencing to collect the gas in the jars or in a 
gasometer, a certain quantity, not too small, must be allowed to 
escape into the air, until there is no doubt that all the air which 
was contained in the WoulfFs bottle has been driven out. If a 
mixture of hydrogen and air is set on fire, the combustion would 
probably be accompanied by a powerful explosion, and the vessel 
shattered. 

In its chemical behaviour, hydrogen is particularly distinguished 








20 


Text-Book of Inorganic Chemistry . 

from oxygen by the fact that it is combustible ; it takes fire when 
ignited in contact with air, and burns with a scarcely luminous 
flame. 

The chemical properties of hydrogen generally are best seen in 
its behaviour to oxygen, in connection with which we have to con¬ 
sider four distinct questions :— . . 

i. Under what conditions does the chemical combination of 

hydrogen and oxygen take place ? 

ii. What is the product of this union ? 

iii. With what phenomena is the chemical union of the two 
substances accompanied ? 

iv. What regularities do we observe in the process ? 

Oxygen and hydrogen gases, when mixed together in any pro¬ 
portion, remain chemically unchanged under ordinary circum¬ 
stances for any length of time. Direct sunlight does not act upon 
the mixture, as it does on the mixture of other gases. But if such 
a mixture be raised to a certain temperature, for which the passage 
of an electric spark or contact with a burning chip of wood suffices, 
chemical union at once follows, in this case with an explosion. In 
order that the hydrogen may unite with the oxygen, it must be 
raised up to or above a certain temperature, called its temperature 
of ignition. 

The product of the combination of hydrogen and oxygen is, 
under all conditions, water. Of this we may easily convince our¬ 
selves by a simple experiment. 

Although it is dangerous to ignite a mixture of hydrogen and 
oxygen, a burning body may be safely brought near a jet of 
hydrogen gas issuing from a small opening— e.g. a glass tube into 
the air. The gas then catches fire and burns with a pale flame. 
If a tube, bent upwards at its point, is connected with a hydrogen 
apparatus or a gasometer containing hydrogen, the gas set on fire, 
and the tube then dipped into a large glass flask containing dry 
oxygen, the flame of the hydrogen becomes smaller—and is at the 
same time tinged yellow—and the inner walls of the flask become 
covered with drops of dew. The longer the hydrogen burns in 
the oxygen the more does the flask become bedewed, and the 
greater is the quantity of water produced. 

The same phenomenon is perceived when the flask is simply 
filled with atmospheric air, the oxygen of which then serves to 
convert the hydrogen into water. 


Hydrogen . 21 

If in this experiment oxygen is led into the flask at the same 
rate as it is consumed by the hydrogen, and if care is taken to 
keep down the high temperature produced by the combustion by 
cooling the outside of the flask, a considerable quantity of water 
may in time be produced from its constituents. 

The chemical union of hydrogen and oxygen, whether con¬ 
tinuous by burning the gas in air, or whether instantaneous by the 
explosion of the two gases previously mixed, is accompanied by 
a very large evolution of heat. In daily life the temperature of a 
flame is often judged by the light which it emits, and it is cus¬ 
tomary to consider the flame of a spirit-lamp less hot than that ot 
a candle. The chemist knows that this measure of the tempera¬ 
ture of different flames is wrong, and that the scarcely luminous 
flame of hydrogen burning in oxygen possesses one of the highest 
temperatures which we can produce by the processes of com¬ 
bustion. Platinum, which is so difficult to melt, and which re¬ 
mains unchanged in the strongest heat of our ordinary furnaces, 
may be easily fused in the oxy-hydrogen flame. 

The explosive action produced by the ignition of a mixture of 
hydrogen and oxygen in the proper proportions and the danger 
attending these explosions make it necessary to be careful in 
experimenting with such a mixture of the two gases. If a large 
glass jar standing over water were filled two-thirds full with 
hydrogen and the other third with oxygen, the mouth closed with 
a glass plate, the jar inverted, and, at the moment when the plate 
was taken off, a burning taper applied, not only would the mixture 
burn with a loud explosion, but the glass cylinder would probably 
be broken. 

It would be still more dangerous to ignite a gasometer lull of 
the mixed gases by opening the upper stop-cock and applying a 
flame. Since the invisible particles of hydrogen and oxygen are 
most intimately mixed with one another, certainly much closer 
than we can bring the particles of two solid substances by rubbing 
them together, the ignition at one point of the gas would be imme¬ 
diately transmitted through the entire mass, and a loud explosion 
with the destruction of the vessel, would be the result. 

In order to exhibit this phenomenon in a harmless manner, the 
mixed gases must be enclosed in very thin membranes. A thin 
calf’s bladder may be taken, but it is better and even less dan¬ 
gerous to explode the mixed gases in soap bubbles. The bubble 
filled with the gases may be prepared in the following way. 


22 Text-Book of Inorganic Chemistry. 

About two volumes of hydrogen and one of oxygen are mixed 
in a glass bell-jar (c, fig. 8), standing over water, which is provided 
with a brass stop-cock. To this brass stop-cock is screwed on 
another stop-cock, to which is fastened a bladder softened in warm 
water and squeezed together to expel the air. 
If now both taps are opened and the bell- 
jar pressed down in the water, the mixed 
gases are forced up into the bladder. As 
soon as the bladder is filled the taps are 
again closed and the bladder unscrewed. 
On dipping then the end of a tube attached 
to the brass stop-cock of the bladder under a 
little soap and water in a saucer, opening 
the cock, and pressing the bladder gently, a 
large number of soap bubbles, filled with the 
mixed gases, may be obtained on the saucer. 
If then a burning taper is brought into con¬ 
tact with these soap bubbles, a violent ex¬ 
plosion is produced, but without in the least 
damaging the saucer. The bubbles may even be exploded in the 
hollow of the hand without the least danger : no shock is expe¬ 
rienced, but only a feeling of gentle warmth. 

If a thin glass flask is filled with the explosive material, the 
mouth closed with a cork, and well wrapped in towels so that the 
mouth just projects, then on pointing the mouth towards a bare 
wall, withdrawing the cork and applying a flame, a smart explosion 
results, and the glass flask is burst into a thousand pieces, which 
remain in the towels. 

A similar, though much less powerful action, is produced by the 
ignition of a mixture of two volumes of hydrogen and five volumes 
of common air. 

In all cases the explosion is caused by the high temperature 
produced when a mixture of hydrogen and oxygen is burnt. The 
water vapour which is produced, and which is momentarily heated 
to redness, expands so largely and so suddenly that thin vessels 
cannot stand the pressure and are therefore burst. Immediately 
after this expansion follows the contraction and condensation of the 
hot water vapour to liquid water by the cooling which it suffers on 
coming into contact with the surrounding cold air. By this means 
a partially vacuous space is formed, and the rushing together of the 
surrounding air, together with the previous expansion, produces the 



Fig. 8. 









Hydrogen. 


23 


movement in the air which is perceived by the ear as the sound of 
the explosion. 

If such expansions and contractions follow one another alter¬ 
nately, regularly, and quickly, they may set a column of air in 
vibration, and produce musical notes. This may be easily shown 
by lowering an open glass tube, of about a metre long, vertically 
over a flame of hydrogen gas, so that the flame continues to burn 



(fig. 9). The cold air rises rapidly in the tube, mixes with the 
hydrogen gas, and causes it to burn more energetically. And since 
the red-hot water vapour which is continually produced is as con¬ 
tinually cooled by the ascending air, small vacuous spaces aie 
formed, and the whole column of air is set into vibration. A 
musical note is therefore heard, the pitch of which depends partly 































24 


Text-Book of Inorganic Chemistry. 

on the width and length of the tube, and partly on the position of 
the flame burning in it. In consequence of these vibrations the 
flame itself is set into quivering motion, something like fig. io. 

Several practical applications are made of the high temperature 
produced by the combustion of a mixture of hydrogen and oxygen. 

Metals which do not 
melt at the highest tem¬ 
perature of our furnaces 
— e ‘g platinum—may be 
easily fused by means of 
the oxy-hydrogen blow¬ 
pipe. This blowpipe 
(fig. ii) is so arranged 
that the hydrogen and 
oxygen only mix at the 
moment when, and at 
the place where, they 
issue from the blowpipe. 

It consists of a copper 
tube, which opens into 
one of platinum, and into 
which hydrogen is led by 
the tap H. A second 
narrower tube of copper 
is placed inside the > 
larger one, and also ter¬ 
minates in a platinum point. Oxygen is led into this tube by the 
tap O, and it can be raised or lowered at will. 

As crucible for the molten metal a block of quick-lime is used. 
This block is sawn through and hollowed out as shown in the 
figure. The lower piece B, which is provided with a spout D, 
receives the metal to be melted ; the upper piece A is pierced at 
the top to receive the end of the blowpipe. 

The highest temperature is obtained when the taps O and H 
are so placed that for one volume of oxygen two volumes of hydro¬ 
gen are burnt, which may be easily seen from the nature of the 
flame after a little practice. The molten platinum is poured from 
the opening D. 

Iron, copper, and other metals may also be easily melted in 
the same manner ; but the fusion of those metals which become 









Hydrogen. 2 5 

oxidized in the presence of oxygen at a high temperature must be 
effected with an excess of hydrogen. 

Lead, which belongs to the most easily fusible of the heavy 
metals, cannot be melted by a small flame, when in large plates, 
owing to its high conducting power for heat. In' order to melt 
together air-tight the immense sheets of lead of which the 
chambers of the sulphuric acid works are constructed, the oxy- 
hydrogen blowpipe is successfully employed. Only the high tem¬ 
perature of this flame can suffice to melt together the edges of two 
thick plates of lead. 

The high temperature of a burning mixture of hydrogen and 
oxygen is also used for the production of a brilliant light. The 
mixed gases when ignited produce a hardly luminous flame ; but 
this flame can make solid infusible substances brought into it so 
hot that they give out the most intense light. Quicklime is par¬ 
ticularly adapted for this purpose ; it does not melt in the flame, 
but remains unchanged. The light which a piece of lime emits 
when it is brought into the oxy-hydrogen flame, with suitable 
arrangements, is so intense that it is used to illuminate large 
spaces at night, and is well adapted for signals— e.g. lighthouses. 
This source of light, which is also used for magic-lanterns and 
other purposes, is called the lime-light , or sometimes the Drum¬ 
mond-light, after Drummond, who first brought it into use. 

The fourth of the questions previously asked (p. 20) remains 
unanswered: What regularities do we observe on the union of 
hydrogen and oxygen to form water ? 

We have already learnt (p. 6) that on the electrolysis of water 
with two platinum plates as electrodes exactly two volumes of 
hydrogen are set free for every one volume of oxygen. This leads 
to the conclusion that, on the synthesis of water, hydrogen and 
oxygen would unite in the same proportion. Experiment has, in 
fact, proved that if one volume of oxygen unites chemically with 
hydrogen, exactly two volumes of the latter gas are required, and 
that when an excess of either of the gases is employed, this excess 
remains behind unchanged. This regularity may be illustrated in 

the following manner. . 

An eudiometer of simple construction and sufficient for this ex¬ 
periment (fig. 12), is prepared from a glass tube of i£ centimetre 
internal diameter and 50 centimetres length. The upper end 
is closed, and two stout platinum wires are melted into it, of which 


2 6 


Text-Book of Inorganic Chemistry. 


the ends are about 5 mm. from one another, and are connected 
together by a thin platinum wire. In order to calibrate this tube 
so that a number of equal volumes of gas may be marked outside, 


it is filled with water, inverted, 
and the air from a small mea¬ 
suring tube, about 5 mm. in 
diameter and 6 cm. in length, 
allowed to ascend into it. The 
eudiometer is next lowered in 
the cylinder (fig. 12) until the 
level of the water inside and 
outside the tube is the same, 
and then an india-rubber ring 
is slipped along the tube to 
mark the level of the water. 
After repeating this six or 
eight times, as many divisions 
of the tube are obtained, 
marked externally by the in¬ 
dia-rubber rings, which need 
not be moved from their po¬ 
sitions if the instrument is 
carefully handled. 



By means of this eudio¬ 
meter it may be easily shown 


that when equal volumes of hydrogen and oxygen are introduced 
into it and raised to their temperature of ignition, one half of the 
oxygen remains behind unchanged, while the other half combines 
with all the hydrogen to form water. 

The small measuring tube is filled with oxygen from a small 
flask containing the gas, which is closed with a cork and stands 
over water, and the gas is passed up into the eudiometer. The 
operation is repeated, and then two tubes full of hydrogen are 
passed up (fig. 13). In order to be quite sure that the mixture 
reaches to the fourth ring on the tube, the latter is transferred to 
the cylinder and lowered until the level of the water is the same in¬ 
side and outside the tube. 

The eudiometer is now closed with the thumb, and transferred 
to a small pan with water, for which an ordinary porcelain or glass 
mortar may be used, and firmly pressed with the hand on a thick 
piece of india-rubber at the bottom in order to prevent expulsion 









27 


Hydrogen. 

)f the gas on the explosion of the mixture (fig. 14). The ignition 
s brought about by connecting the two platinum wires at the end 



Fig. 14. 

of the eudiometer with the wires from a battery of four or six 
Bunsen’s cells. The electric current then heats the thin platinum 
























28 


Text-Book of Inorganic Chemistry. 

so strongly that it reaches the temperature of ignition of the mixed 
gases. 1 The combination is shown by a pale flash of light and a 
feeble, scarcely perceptible shaking of the tube. 

That a condensation has taken place owing to the union of the 
two gases to form liquid water, or, in other words, that a diminution 
of the volume of the gases has been produced, only becomes visi¬ 
ble when the eudiometer is slightly inclined and the india-rubber 
plate removed. The water then rises in the eudiometer, and when 
dipped in the cylinder reaches exactly to the first mark. 

Of the four volumes of gases three have, therefore, apparently 
disappeared. That the small remaining quantity of gas is oxygen 
may be easily seen by closing the eudiometer with the thumb ; in¬ 
verting and plunging a glowing chip quickly into the gas, the chip 
at once catches fire. 

If an alteration is made in the experiment by taking hydrogen 
in excess—for example, by taking three volumes of hydrogen and 
one volume of oxygen—of these four volumes one volume would 
also remain behind after the explosion, and would now be hydrogen. 
That the remaining gas is really hydrogen may be readily proved 
by its inflammability. 

As will be easily understood, such experiments do not give 
exact nor even approximately exact results. The water over which 
the measurements are made always contains air which mixes with 
the other gases, and in apparatus constructed and managed in this 
way the influence of temperature and of pressure cannot be suffi- 
s ciently taken into account. But the experiments suffice to show 
all they are intended to—viz. the proportion by volume in which 
hydrogen and oxygen unite to form water. 

A further question which must also be mentioned here is, 
What space does the water gas occupy which is produced by the 
union of two volumes of hydrogen with one volume of oxygen, and 
in what proportion does the volume of the water gas stand to that 
of its components ? Since water under the ordinary pressure of 
the air only becomes completely gaseous at ioo°, the volume of 
the water gas must be measured at some temperature above ioo°. 
Experiment teaches us that the volume of water gas, measured at 

1 Instead of connecting the two wires in the eudiometer with a thin one of 
platinum, we can leave the ends free without touching one another inside the 
tube. If we then connect the wires from the battery with an instrument called 
an induction coil and the wires from the coil with those of the eudiometer, 
sparks pass between these wires w'hen the connections are completed.— Ed. 


29 


Hydrogen. 

a temperature of about 150°, which is produced from two volumes 
of hydrogen and one volume of oxygen also at the same tempera¬ 
ture does not occupy three volumes as we might expect, but only 
two volumes. It follows, therefore, that when hydrogen and oxygen 
unite to form water gas a condensation of the volume takes place 
in the proportion of three to two. 

Since we know the relative densities of hydrogen and oxygen, 
in what proportion by volume they combine with one another, and 
what condensation is produced on combination, we can easily find 
the relative density of water gas. 

It has been proved that two volumes of hydrogen (weighing 
2 x 0-0692) combine with one volume of oxygen (weighing 1-1056) 
to form two volumes of water gas, whence— 

2 vols. hydrogen weigh 2 x 0-0692 . - 0*1384 

1 „ oxygen .......= 1-1056 

2 „ water gas ...... 1*2440 

and, therefore* one volume of water gas weighs 1 2 ^ 4 ° = 0-622, a 

number agreeing almost exactly with the experimental results of its 
specific gravity. 


COMBUSTION. 

We have seen that a number of bodies—sulphur, charcoal 
phosphorus, iron, hydrogen, &c.—when heated to their tempera¬ 
ture of ignition, burn in oxygen, and that the products are always 
compounds of the burnt bodies with oxygen. 

It may be asked, Is oxygen the only gas in which combustible 
substances can burn ? or Do other bodies also possess this property? 
Experiment has long answered this question. 

Among other bodies which possess this property , is the element 
chlorine—a greenish-yellow gas—which, together with the metal 
sodium, makes up common salt, and which may be easily obtained 
from this substance. If hydrogen is allowed to stream out of a 
small opening in a glass tube, is ignited in the air, and then 
plunged into a jar of chlorine, the gas goes on burning, and the 
yellow colour of the chlorine gradually disappears. The flame of 





30 Text-Book of Inorganic Chemistry. 

the hydrogen burning in chlorine is not yellowish as when it burns 
in oxygen, but of a greyish colour, and the product is not water but 
a gaseous compound of hydrogen and chlorine, possessing strongly 
acid properties, and called hydrochloric acid. 

This one experiment proves, and it might be confirmed by 
hundreds of others, that other substances as well as oxygen can 
support the combustion of burning bodies. 

The process of combustion in oxygen, which is by far the com¬ 
monest, is distinguished from combustion in other gases, and is 
called oxidatio 7 i. Sulphur, it is said, becomes oxidized or suffers 
oxidation when it burns in oxygen, or generally when it unites 
chemically with this substance. 

Combustion, and particularly that of wood, belongs to the phe¬ 
nomena of nature which were earliest noticed by man. The 
knowledge of this process is as old as the observation that the 
heavenly bodies move—movements which were, however, recog¬ 
nized and their investigation attempted thousands of years before 
it was known that the earth moves round the sun. And the process 
of combustion remained a secret to investigators even later. It was 
a century and a half after Galileo’s words, ‘ e pur si muove,’ before 
this process was first correctly explained by Lavoisier. The -neces¬ 
sary prelude was the discovery of oxygen, which was closely 
followed by Lavoisier’s theory of combustion. 

The history of the theories of combustion teaches us, if anything 
can, how mere philosophizing on natural phenomena, without an 
experimental basis, is a vain and useless undertaking. 

At the end of the seventeenth century, the celebrated German 
chemist Stahl put an end to the obscure and contradictory ideas of 
earlier centuries by his phlogiston theory. The known fact that wood 
when burnt on the hearth gradually disappears, and that something 
apparently leaves it. with- the flame, for only ashes remain behind', 
led Stahl to the hypothesis that wood and all combustible bodies 
must contain a volatile substance, which on their combustion is 
given off with the production of heat and often of light. This 
substance he called phlogiston , and the process of combustion 
dephlogistication. Sulphur, phosphorus, or charcoal, when burnt, 
were said to be dephlogisticated or deprived of their phlogiston, 
and the solid product of the combustion of iron was dephiogisti- 
cated iron. 

When metallic iron was obtained by heating dephlogisticated 
iron with charcoal, it was said that the dephlogisticated iron 


Combustion. 


3i 


became phlogisticated, or provided with phlogiston. Iron would 
accordingly be a more complex substance than the product 
obtained on burning it. 

This hypothesis—the phlogiston theory —imparted such a 
simple and satisfactory explanation to all the then known pheno¬ 
mena of combustion, that for more than half a century no one 
doubted its correctness. On the other hand, Stahl, as well as 
some of his contemporaries, had made observations which strongly 
contradicted the theory ; but no attention was then bestowed on 
them, and their importance was overlooked. 

It had been long observed by various chemists, and was not 
discovered by Lavoisier, that many metals on dephlogistication— 
i.e. on burning—increased in \yeight, notwithstanding that they 
were supposed to lose phlogiston. That this loss of phlogiston 
ought not to produce an increase in weight, but a diminution, was 
however, considered of no importance. 

Only after the discovery of oxygen by Priestley and by Scheele 
did Lavoisier prove, by numerous new determinations, that bodies 
on combustion simply combine with oxygen, and that the increase 
in weight is the same as that of the oxygen consumed. 

This result, now so familiar to us, may be easily proved by a 
simple experiment. A small quantity of iron in the state of a fine 
powder is placed in a bulb-tube, and then accurately weighed. A 
stream of oxygen is now allowed to pass through the bulb-tube, 
and the iron is at the same time heated ; suddenly the latter 
begins to glow, and this heating effect is transmitted through the 
entire mass in consequence of the union of the iron and oxygen. 
If then the bulb-tube is again weighed, when perfectly cold, it will 
be found to have considerably increased in weight. The same 
method may also be employed in order to determine quantitatively 
how much oxygen is consumed by a known weight of iron. 

It may be here remarked that although iron may increase in 
weight on combustion in oxygen, a piece of wood or a candle 
clearly diminishes in volume and weight when burnt. But in these 
cases appearances are deceptive. Chemists have determined not 
only that the substance of wax-candles increases in weight on 
burning, but also that the products of combustion weigh more than 
four times as much as the wax consumed. This is not ordinarily 
perceived, as these products of combustion are volatile and 
invisible to the eye. 

That even a candle increases in weight on burning, or that the 


3 2 




Text-Book of Inorganic Chemistry. 


products of combustion weigh more than the burnt wax, may 
easily be made visible by the apparatus shown in fig. 15. 



Fig. 15. 


A glass cylinder (an ordinary Argand lamp cylinder), a , is placed 
above a cork, pierced with several holes, upon which a piece of a 



































Combustion. 


33 


wax-candle, about two inches long, is fastened. The candle is sur¬ 
rounded with a small piece of tinfoil, which receives any molten 
wax running down. 

The upper end of the cylinder is connected with a bent wide 
glass tube b, by means of an india-rubber stopper, and made to fit 
air-tight. In order to protect the stopper from the heated air 
rising from the flame, two perforated pieces of platinum foil are 
placed in the tube beneath it, care being taken that the perforations 
do not coincide with one another. 

The vessels £, c , d, e, which communicate with the cylinder a , 
serve partly to retain the products of combustion of the wax, which 
consists of carbon, hydrogen, and a small quantity of oxygen, and 
partly to make these products visible to the eye. The tube b is 
empty, and receives a large portion of the water produced; the 
small flask c contains clear lime-water, which becomes milky when 
carbonic acid, passes through it owing to the production of calcium 
carbonate ; the tubes d and e are filled with pieces of caustic soda, 
which retain the remainder of the water and carbonic acid. A 
constant stream of air, which serves both to make the candle burn 
and.to pass the products of combustion through the vessels b , c, d , 
e , is maintained by the Bunsen’s pump P, which is in communica¬ 
tion with these vessels by the india-rubber tubing supported at g. 

The various portions of the apparatus are fastened to the glass- 
rod z, z, which is attached to one arm of an ordinary pair of scales, 
and is then exactly balanced by weights placed in the pan attached 
to the other arm. On opening the tap of the pump, no change 
takes place in the equilibrium. The candle is now carefully with¬ 
drawn, ignited, and again placed in the tube. 

After a short time the inner surface of the tube b becomes 
covered with moisture, and at the same time the lime-water in c 
becomes turbid. And as the combustion proceeds, drops of water 
collect in the tube b, and the arm of the balance to which the ap¬ 
paratus is attached gradually sinks, until it rests on the foot of the 
instrument. 

If the process is continued long enough to entirely consume the 
candle, a considerable weight must be placed in the pan of the 
balance in order to again restore equilibrium, which is sufficient 
proof that the substance of the candle, the wax, when burnt to car¬ 
bonic acid and water, increases in weight. 

We are so accustomed to see the burning of wood, oil, sulphur, 

D 



34 Text-Book of Inorganic Chemistry. 

phosphorus, and other combustible bodies in the free atmosphere, 
that we are inclined at first to imagine that the burning bodies 
necessarily require a surrounding atmosphere rich in oxygen. Ihis 
idea is also erroneous. If we drop a piece of phosphorus into a 
vessel of warm water, so that it melts, and then lead oxygen gas to 
it under the water in a slow stream, the phosphorus burns under 
the water, producing a considerable quantity of light and heat. The 
product is the same as when phosphorus burns in free air—viz. 
phosphoric acid. 

For the conversion of phosphorus into phosphoric acid, the 
presence of free gaseous oxygen is not even a necessary condition. 
In nitric acid we are acquainted with a liquid chemical compound 
of nitrogen, rich in oxygen, which can be easily decomposed. If 
we heat a small piece of phosphorus in nitric acid, a portion of the 
oxygen of the latter body combines with the phosphorus without 
first assuming the gaseous state, and finally entirely oxidizes the 
phosphorus to phosphoric acid. 

Metallic tin— eg. tinfoil—which, when heated in the air or in 
oxygen, is converted into a white substance (stannic oxide, or putty 
powder), suffers the same change when acted on by nitric acid. 

Charcoal may also be easily burnt by nitric acid to form car¬ 
bonic acid. If a small piece of charcoal is made red-hot at one 
end, and dipped into a flask containing fuming nitric acid, so that 
the glowing point just touches the liquid, energetic combustion 
takes place, with a large evolution of light and heat. 

In the same manner as free oxygen gas and the loosely united 
oxygen of nitric acid can oxidize combustible bodies, solid sub¬ 
stances which contain their oxygen or a portion of it loosely com¬ 
bined may also be used for the same purpose. Powdered nitre— i.e. 
potassium nitrate—when sprinkled on a glowing coal, causes the latter 
to burn brilliantly to carbonic acid. Potassium chlorate acts in the 
same manner, and not only towards charcoal, but also towards 
other combustible bodies. If a small quantity of dry, powdered 
potassium chlorate is rubbed in a mortar with powdered sulphur 
(only very small quantities must be used), the heat produced by 
the rubbing is sufficient to cause a number of small explosions, 
depending upon the oxidation of the sulphur by a portion of the 
oxygen contained in the potassium chlorate ; explosive compounds 
of chlorine containing less, oxygen are also produced at the same 
time. 

Ordinary combustions in oxygen are usually accompanied with 


Combustion. 


35 


an evolution of light, and heat is always set free. In the case of 
some combustions, however, no light appears, and these are usually 
accompanied by less sensible heat. These processes of oxidation 
may be called slow combustion, in contradistinction to those pro- 
uang ight and sensible heat, called quick combustion. Phos¬ 
phorus, which when heated in the air to its temperature of ignition 
burns quickly and brilliantly to phosphoric acid, also becomes oxi¬ 
dized at the ordinary temperature in the air, without any light being 
perceptible in the daytime. The product of oxidation in this case is 
an oxide containing less oxygen than phosphoric acid, and is called 
phosphorous acid. A piece of phosphorus, which when heated in 
the air undergoes quick oxidation and is converted into phosphoric 
acid in a few minutes, requires as many months, or even longer, 
to be completely converted into phosphorous acid by slow com¬ 
bustion. 


The continuous oxidation produced in the human body by 
inspired oxygen belongs to those processes of slow combustion 
which are unaccompanied by any evolution of light. The oxygen 
brought into the lungs by inspiration is taken up by the arterial 
blood, and is so distributed to various parts of the body. It then 
oxidizes the various tissues of the body with which the blobd comes 
in contact, as well as portions of the blood itself, being converted 
into water and carbonic acid. The latter substance is chiefly taken 
up by the venous blood and discharged from the lungs in the 
expired air. That the expired air is rich in carbonic acid may be 
easily shown by blowing through lime-water with a piece of glass 
tubing; the clear liquid becomes milky, and deposits a white 
sediment of calcium carbonate. 


We know that when hydrogen is allowed to stream from a jet 
into the air, and a burning body brought near, it catches fire and 
burns ; but we might in vain attempt to ignite oxygen under the 
same conditions. Suppose, however, the earth were surrounded 
with an atmosphere of hydrogen, instead of with one of common 
air, and that its inhabitants required hydrogen gas just as we do 
oxygen, how would oxygen behave if a jet of it were heated in 
this atmosphere of hydrogen ? Would it then also bum with a 
flame in the same manner as hydrogen in oxygen ? 

Experiment has answered this question in the affirmative. It 
may be easily shown that oxygen bums in hydrogen, just as hydro¬ 
gen in oxygen. If oxygen is allowed to flow slowly from a gas- 


36 Text-Book of Inorganic Chemistry. 

ometer out of a glass tube bent upwards, and a large inverted jar of 
hydrogen which has been previously ignited is placed over the jet, 
it will be seen that the oxygen catches fire and burns in the atmo¬ 
sphere of hydrogen with a pale flame, until all the latter gas is 
consumed. 

From this it is clear that the distinction between combustible 
bodies and supporters of combustion is not strictly scientific, and 
is as inexact as when we speak of the rising and setting of the sun. 
We must consider the process of combustion in the wide meaning 
of the word as the chemical union of heterogenous bodies, and in 
the narrower meaning as an oxidation process, or the union of a 
substance with oxygen. 

The oxygen may be more or less easily again separated from 
the oxidized bodies. It is only necessary to heat the red oxide of 
mercury, oxide of silver, or the oxides of the other noble metals, in 
order to drive out the oxygen and to obtain the metal again. If 
we pass hydrogen over the product of combustion of copper (black 
oxide of copper), and apply heat, the gas unites with the oxygen 
to form water, and metallic copper remains behind as a red powder. 
The powerful affinity which carbon has for oxygen at a red heat 
makes it well adapted for abstracting the oxygen from the oxides 
of those substances (eg. phosphorus and iron), which possess strong 
affinities for oxygen. By means of glowing charcoal we can reduce 
phosphoric acid to phosphorus, and oxide of iron— eg. the iron-ores 
employed in blast furnaces—to metallic iron. 

By the processes of reduction we produce exactly the opposite 
effect to those of combustion, or more correctly of oxidation 
generally k We understand by reduction not only the production of 
the elements from their oxygen compounds, but also their sepa¬ 
ration from compounds with other substances. We reduce chloride 
of silver by separating the silver in some suitable manner from 
chlorine ; we reduce mercury from cinnabar (mercuric sulphide) by 
the removal of the sulphur, &c. 

Other reactions are also called processes of reduction, by which 
compounds of oxygen, sulphur, chlorine, &c., are converted into 
other compounds containing less of these last-named elements. 
We reduce the salts of ferric oxide to those of ferrous oxide, which 
contain less oxygen, cupric chloride to cuprous chloride, containing 
less chlorine, sulphuric acid to sulphurous acid. 

In the same manner oxidation means not only the combination 


The Laws of Chemical Combination. 37 

of a substance with oxygen, but also the addition of more oxygen, 
sulphur, chlorine, &c., to compounds already containing these 
elements. We oxidize the salts of ferrous oxide to those of ferric 
oxide, cuprous chloride to cupric chloride, sulphurous acid to sul¬ 
phuric acid, &c. 


THE LAWS OF CHEMICAL COMBINATION. 

We have seen that the union of hydrogen and oxygen to form 
water always takes place according to a definite proportion. Ex¬ 
actly two volumes of hydrogen always combine with exactly one 
volume of oxygen. Now since we know the specific gravity of the 
two gases, or that one volume of oxygen (sp. gr. = rio6) weighs 16 
times as much as one volume of hydrogen (sp. gr. = 0-069), it follows 
that the quantity of oxygen (one volume) which unites with two 
volumes of hydrogen weighs 8 times as much as the hydrogen. 
From this, the percentage composition of water by weight may be 
easily calculated : 

2 vols. hydrogen . . . . =2x0-069=0*138 

1 „ oxygen.=1*106 

1 244 

If, then, x is the weight of hydrogen and y the weight of oxygen 
contained in 100 parts by weight of water, we get the following 
simple proportions : 

1-244 : 0-138 :: 100 : x, 

1-244 • 1,106 • • i°° : y, 

which give x= 11 -11 and/= 88-89, and hence 100 parts of water 
contain : 

Hydrogen.= iru parts. 

Oxygen.= 88-89 „ 

Water.= i°°* 00 » 

This percentage composition of water, calculated from the pro¬ 
portions by volume in which hydrogen and oxygen unite with one 
another and from the specific gravities of the two gases, agrees 
exactly with innumerable analyses and syntheses of water which 





38 


Text-Book of Inorganic Chemistry. 

have been made from time to time. 1 It has been further proved 
that pure water, whatever may be its source, whether obtained 
from ice, snow, or water vapour, or whether prepared artificially 
from its constituents, has always the same composition. 

A chemical compound containing oxygen and hydrogen, like 
water, but in some other proportion, is not water. We are 
acquainted with such a compound, called hydrogen peroxide, to 
which we shall refer later on. This body contains only 5-9 per cent, 
of hydrogen, and 94-1 per cent, of oxygen. 

The composition of all the more accurately known chemical 
compounds has, like that of water, been carefully determined, and 
we know not only of what elements they consist, but also in what 
proportions they contain these elements. 

From these results of analytical chemistry two laws of extreme 
importance in discussing the regularities in the composition of 
chemical compounds have been deduced :— 

i. On the chemical union of two substances their original weight 
remains unchanged. 

ii. Every chemical compound contains the simple substances of 
which it is composed in one , and only one, proportion by weight. 

From this, however, it does not follow that two substances 
which contain the same simple substances in the same proportions 
are therefore identical. Numerous isomeric compounds exist, 
which, although they have the same elementary and percentage 
composition, possess different properties, and are quite different 
bodies. 2 

1 One of the simplest methods for determining the composition of water by 
weight depends upon the fact that red-hot copper oxide is reduced in a stream 
of hydrogen to metallic copper, while its oxygen unites with the hydrogen to 
form water. If we collect this water by some compound which will absorb it 
[e.g. calcium chloride) and-weigh the tube containing this substance before and 
after the experiment, we know the weight of water which has been formed. 
And if we also weigh the copper oxide before and after the experiment we find 
the weight of oxygen which has combined with the hydrogen to produce the 
known weight of water. It is then always found that 100 parts by weight of 
water contain exactly 88-89 parts by weight of oxygen, the remaining ii-ii 
parts being hydrogen. —Ed. 

2 Isomeric compounds, or those possessing the same percentage composi¬ 
tion, may be divided into two classes—viz. thpse which have the same mole¬ 
cular weight, but in which the atoms are differently arranged— e.g. propionic 

c c h ( CH 

acid Ico-OH’ an d methyl acetate jcCFOCH 3 ’ called metameric com¬ 
pounds ; and those of which the molecular weight of the one is some multiple 


The Laws of Chemical Combination. 


39 



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40 


Text-Book of Inorganic Chemistry. 

The knowledge of the composition of the thousands of chemical 
compounds at present known to chemists would have been of little 
service to us had we been compelled to rest satisfied with this 
alone. But this knowledge serves as a foundation for determining 
and establishing the laws in which the elements combine with one 
another, and compounds with compounds. 

The lines I.a, II .a, III.a, in the accompanying table give the 
percentage composition of compounds of hydrogen, copper, lead, 
and thallium with oxygen, sulphur, and chlorine respectively. 
The merest glance at these numbers shows that they are very 
different from one another, and that no apparent regularity exists 
between them. The only regularity which occurs in these three 
lines is that copper unites with less oxygen, sulphur, or chlorine 
than hydrogen, lead with still less, and thallium least of all. This 
is made more manifest in the lines Lb, 11 . b, 111. b, in which are 
Stated the relative weights of hydrogen, copper, lead, and thal¬ 
lium, which unite with ioo parts of oxygen, sulphur, and chlorine 
respectively. But the numbers in line I .b : 12*5, 396*25, 1293*7, 
2550*0, which are in the proportion 1 : 317 : 103*5 : 204, have still 
no simple relation to one another. If, however, we compare this 
proportion with that in which the same four elements unite with 
equal weights of sulphur (100 parts), we see from line II .b that this 
proportion is the same, for, dividing 6*25, 198*12, 646*8, 1275*3 
by 6*25, we get as before 1 : 31*7 : 103*5 : 2 °4 1 and further, the 
ratio in which the four elements unite with chlorine (lines III .b, 
III.c) is also the same. 

Here, then, we at once perceive definite regularities which may 
be formulated in the following statement :— 

The same proportional weights of hydrogen, copper , lead, and 
thallium always unite with one definite weight of oxygen, sulphur, 
or chlorine. 

These elements do not, however, occupy any exceptional posi¬ 
tion, and all that has been said for their combining proportions 
holds also for all the others. We thus obtain the following general 
law :— 

The proportion i?i which two bodies (A and B ) combine with a 
third (C) is also the proportio 7 i in which they unite with all other 
bodies as well as with one another. 

of that of the others— e.g. ethylene (C 2 H 4 ) and propylene (C 5 H 6 ), called polymeric 
compounds. Examples of these compounds are more common in organic than 
in inorganic chemistry.—E d. 


The Lazos of Chemical Combination. 41 

Thus, from the accompanying table, we see that the weight of 
chlorine (35*5 parts) which unites chemically with 1 part of 
hydrogen is the same as that which unites with 317 parts of 
copper, 103-5 parts of lead, or 204 parts of thallium. And, again, 
that the same relative quantities of hydrogen, copper, &c., which 
unite with 35-5 parts of chlorine combine also with 16 parts of 
sulphur or with 8 parts of oxygen. 

We have thus obtained the following numbers for these ele¬ 
ments, which represent the relative weights in which they combine 
with one another and with other elements :— 


Hydrogen 

1 -o 

Oxygen 

8-o 

Sulphur . 

16-0 

Chlorine . 

• 35'5 

Copper 

. 317 

Lead 

. 103-5 

Thallium . 

. 204-0 

Again, suppose we wish to 

discover in what proportions 

chlorine and oxygen unite with 

one another, it is sufficient to 


know the proportions in which they combine with a third body. 
We see from the preceding table that 103-5 parts of lead combine 
with 8 parts of oxygen or 35-5 parts of chlorine. The proportion 
8 of oxygen to 35-5 of chlorine is therefore that in which these two 
elements unite with one another. Hypochlorous anhydride con¬ 
tains exactly 8 parts of oxygen united with 35-5 parts of chlorine. 

In this manner the number representing the quantity of any 
other element which combines with 1 part of hydrogen, 8 parts of 
oxygen, 35-5 parts of chlorine, &c., may be easily obtained. 

If we wish to find this number for the metal sodium, and know 
that the percentage composition of its compound with chlorine 
(common salt) is : 

Sodium. 39 ‘ 3 2 

Chlorine.6cr68 

100 00 

we need then only calculate -the quantity of sodium which would 
combine with 35-5 parts of chlorine. This we find to be 23 from 
the proportion : 6o-68 : 39*3^: :35*5 • and this number, therefore, 
expresses the weight of sodium which combines with 1 part of 
hydrogen, 8 parts of oxygen, 16 of sulphur, &c. 

The relative weights in which bodies combine with one another, 









42 Text-Book of Inorganic Chemistry . 

and in which they displace one another from their compounds, are 
called their equivale?it weights'. 317 parts of copper, or 103*5 
parts of lead always combine with 8 parts of oxygen, 16 of 
sulphur, or 35*5 of chlorine. If in a compound of 357 parts of 
chlorine and 103*5 P ar ts of lead, the lead were displaced by 
some other substance— e.g. copper or hydrogen—the quantity 
of this substance required would not be the same as the lead 
displaced, but 31*7 parts of copper, or 1 part of hydrogen, would 
be necessary. In these compounds x part of hydrogen has 
the same effect as 31*7 parts of copper or 103*5 parts of lead, and 
these elements, therefore, displace one another in their compounds 
not in equal but in equivalent weights. 

In order to displace 35*5 parts of chlorine from its compound 
with 1 part of hydrogen or with 204 parts of thallium by iodine, 
127 parts of the latter body are required. On the other hand, to 
displace the 204 parts of thallium by hydrogen, only 1 part is 
necessary. Finally, if in the compound called copper sulphide, 
consisting of 16 parts of sulphur united with 31*7 parts of copper, 
silver be substituted for copper, it is found that 16 parts of sulphur 
unite with 108 parts of the former metal. 


THE LAW OF MULTIPLE PROPORTIONS. 

The number of chemical compounds would be very limited if 
the elements could only combine with one another in one single 
proportion. Experiment has, however, shown that most of the 
elements can unite-with other elements in more than one propor¬ 
tion, many even in five different proportions, thus producing com¬ 
pounds which are usually very different from one another. 

In 3 uch cases, however, the relative quantities of the two ele¬ 
ments combining together are always some multiple of their 
equivalent weight by small whole numbers. As an example of 
this may be instanced the five compounds of nitrogen and oxygen, 
nitrous oxide, nitric oxide, nitrous anhydride, nitric peroxide, and 
nitric anhydride. 

These compounds contain their constituents in the following 
proportions by weight. 



43 


The Law of Multiple Proportions. 


Nitrogen, 


Oxygen, 

parts 


Parts 


Nitrous oxide 
Nitric oxide 


Nitrous anhydride 
Nitric peroxide . 
Nitric anhydride 


parts 

14 

14 

14 

i .4 

14 


16 

24 

32 

40 


8 



From which it is at once seen that the quantities of oxygen united 
with one equivalent of nitrogen (14 parts) are multiples of its 
equivalent weight by the simple integers 112:3:4:5. 

In the same manner one and the same quantity of iron unites 
with two different quantities of chlorine which are to one another 
in the proportion 2 : 3. 

These and innumerable other results form the experimental 
basis of the Law of Multiple Proportions , which may be expressed 
as follows:— 

The elements unite either in their equivalent weights or in 
simple multiples of their equivalent weights. 

The law likewise holds for compounds as well as elements; 
when the former combine to form more complex compounds they 
do so either in their equivalents or in simple multiples of their 
equivalents. The statement of this law is due to Dalton : it is 
only about seventy years old. Dalton first published his theory in 
detail in the year 1808, at a time when exact chemical investigation 
was largely carried on in England, France, and Sweden, but when 
the leading chemists in Germany, following an erroneous path in 
natural philosophy without any real foundation of fact, long 
wasted their powers in worthless speculation. 

The theories of Dalton and his conclusions derived from the 
law of multiple proportions were, however, more permanent and 
fruitful. 


THE LAW OF VOLUMES. 1 


Not only is there a simple relation between the weights in 
which the elements combine together chemically, but for elements 
in the gaseous state there is also a simple relation between the 

1 The Editor is responsible for this statement of Gay-Lussac’s law. 





44 Text-Book of Inorganic Chemistry. 

volumes of the constituents and of the compound, if gaseous, which 
is produced. 

It will be remembered that when hydrogen gas and oxygen gas 
unite with one another chemically, they always do so in the pro¬ 
portion of two volumes of the former to one volume of the latter, 
and the volume of the water vapour produced (above ioo°) is 
always two volumes. Further, one volume of hydrogen gas always 
unites with exactly one volume of chlorine gas, producing exactly 
two volumes of the gaseous compound called hydrochloric acid. 

These may serve as examples of a general important law—the 
law of volumes—which may be formulated as follows :— 

The volumes of two gases uniting chemically with one another 
are always in some simple proportion , and the volume of the 
compound produced , when gaseous , is also always in some simple 
proportion to those of its constituents. This law also holds for 
compound as well as for elementary gases. We owe its discovery 
to the celebrated French chemist Gay-Lussac, after whom it is 
often called. 


THE ATOMIC THEORY AND CHEMICAL VALENCY 
OR ATOMICITY. 

The stability with which two elements remain combined after 
their chemical union, the impossibility to again recognize in water, 
for example, any property of either of its components, and again to 
break up water by any mechanical means into hydrogen and 
oxygen, led to the idea that on the chemical combination of two 
bodies a penetration of matter and a similar coalescence occurred 
as that when a mixture of copper and zinc are melted together to 
form brass. In such alloys neither of the constituents can be dis¬ 
tinguished or separated by mechanical means. It is impossible, 
even with the most powerful microscope, to perceive the copper 
and zinc in brass, or the gold in gold amalgam. Still, these alloys 
are not generally considered as chemical compounds, because the 
metals of which they are composed generally unite in almost any 
proportion, while the combination of bodies to form real chemical 
compounds always takes place in certain definite proportions. 

If we melt together copper and zinc in those proportions by 
weight in which they both unite with oxygen, sulphur, chlorine, &c. 



The Atomic Theory. 45 

—i.e. in their equivalent weights, an alloy is produced which might 
be considered as a chemical compound ; but such an alloy is also 
produced when we employ a little more or a little less copper or 
zinc than exactly corresponds to their equivalent weights. In no 
case does either copper or zinc remain behind unchanged, so that 
we can separate it from the compound. 

But it is otherwise with real chemical compounds. As the 
metals may be fused together in any proportion, so may hydrogen 
and oxygen be intimately mixed in all proportions. But if we 
transform this mechanical mixture of the two gases into their 
chemical compound—viz. water, we know that the smallest 
quantity of hydrogen or oxygen present in the mixture above the 
equivalent proportion remains behind uncombined and unchanged. 

If the chemical union of two substances depended upon an 
intimate coalescence and penetration of the two bodies, it would 
be difficult to understand why combination only occurs in certain 
definite proportions by weight, and it would be incomprehensible 
that, when it does occur in different proportions, this only happens 
per saliiim , in simple multiple proportions. 

Dalton’s atomic theory affords a simple and satisfactory ex¬ 
planation of these results. 

We believe, as Dalton did, that the substances which compose 
chemical compounds are not capable of indefinite division, but 
that they consist of particles, which cannot be further divided, 
either by physical or by chemical means. These particles are 
therefore called atoms. We imagine that chemical combination is 
produced by the attraction and juxtaposition of the atoms. The 
more firmly the atoms are united together, the stronger must be 
the force by which the combination is produced and maintained. 
This chemical force of attraction is quite different from any other 
physical force. It differs, for example, from the force of gravity, 
which acts through great distances, by which the earth is attracted 
to the sun, and a falling stone to the earth ; it differs also from the 
forces of cohesion and adhesion. Chemical attraction never acts 
at a distance, like gravity, but requires immediate contact of the 
two bodies, just as the passage of an electric current requires the 
direct contact of the conducting bodies. Chemical attraction, as 
we have seen (p. 5), differs also from cohesion, by producing a 
material change in the bodies upon which it acts. 

The chemical force which resides in the atoms, called chemical 
affinity or attraction , is evidently related to the force of electricity. 


46 Text-Book of Inorganic Chemistry. 

Many chemical processes are produced by electricity—combina¬ 
tions as well as decompositions. In the electrolysis of water, the 
hydrogen atoms are set free at the negative pole, and the oxygen 
atoms at the positive pole. And since bodies charged with oppo¬ 
site kinds of electricity attract one another, the conclusion might 
be drawn that the hydrogen which passes to the negative pole is 
charged with positive electricity and the oxygen with negative. 

In the same manner, on the electrolysis of metallic salts, the 
metals are liberated at the negative pole, and the acids or haloid 
constituents at the positive pole. 

From these and other similar results, it was thought that the 
atoms were charged in different degrees with positive or negative 
electricity, and that the combination of two heterogeneous atoms 
was caused by the attraction of the different electricities : this 
was the fundamental idea in the electro-chemical theory. 

But this theory leaves so many phenomena unexplained, and 
even contradicts many chemical reactions, that we do not at pre¬ 
sent employ it, although the existence of near relations between 
chemical and electrical forces cannot be denied. The fact which 
particularly contradicts the theory is, that the atoms of one and 
the same element, which were supposed by the theory to be 
charged with the same kind of electricity, attract one another and 
one generally form very stable compounds. 

We content ourselves with the hypothesis that a force produc¬ 
ing chemical phenomena resides in the elementary atoms, and 
that the atoms of the elements, according to their nature, offer 
one or more points of attraction to other atoms. 

When the atoms of heterogeneous elements attract one another 
by the forces present in them, and form a chemical compound, the 
number of the atoms of the one element which can be attracted 
by one atom of the other element is always limited, and, as we may 
imagine, depends upon the number of points of attraction which 
the atoms possess. This number is different with different elements. 

We know that when hydrogen and chlorine unite with one 
another chemically they always do so in one definite proportion. 
From this and other results we conclude that the atom of hydrogen 
only offers one point of attraction to that of chlorine and' to those 
of other elements. 

The result that when hydrogen unites with oxygen to form 
water twice as much of the former gas is required as when it 
unites with chlorine leads to the conclusion that the atom of oxygen 


The Atomic Theory. 47 

possesses two points of attraction with which it secures the two 
atoms of hydrogen. 

When carbon unites with hydrogen to form marsh gas and with 
oxygen to form carbonic acid, the carbon atom does not fix the same 
number of atoms of the two substances. The atom of carbon, as 
proved from the composition of innumerable compounds, possesses 
four points of attraction for the atoms of other elements, and 
therefore requires four atoms of hydrogen for saturation and only 
two of oxygen, which offer, not one, but two points of attraction to 
other atoms. 

Two atoms of oxygen thus play the same part, in combination 
with one atom of carbon and with all other elements, as four 
atoms of hydrogen ; or, what is the same thing, one atom of 
oxygen has the same value or valency as two atoms of hydrogen. 

Those elements of which the atoms, like those of hydrogen, 
only possess one point of attraction are called monads (mono¬ 
valent or monatomic), those with two points of attraction dyads 
(divalent or diatomic), those with three triads (trivalent or tri- 
atomic), those with four tetrads (tetravalent or tetratomic), &c. 
No elementary atom appears to possess more than seven or eight 
points of attraction, and, as we have remarked, the number of 
atoms of an element which can be fixed by one elementary atom 
is, therefore, limited. 

Although there are some organic compounds which contain 
several dozen atoms united with one another, we know how to 
resolve these apparently complicated proportions into others more 
simple, resembling the inorganic compounds. 

The power of the’elements to fix a certain number of the atoms 
of other elements is called their atomicity or valency , and we say 
that the atomicity of hydrogen is 07 ie , that of oxygen two , and that 
of carbon two or four. Every element possesses a definite 
greatest atomicity, or, in other words, a certain maximum number 
of points of attraction of its atom for other atoms. But the atoms 
of those elements possessing higher atomicity do not always exert 
their full power. Sulphur, of which the maximum atomicity is six, 
and which occurs as a hexad element in sulphuric acid, plays the 
part of a tetrad in sulphurous acid, and of a dyad in sulphuretted 
hydrogen. 

We imagine that when all the points of attraction of an atom 
are not active they are in a state of rest, or that they are latent, 
until aroused by some new chemical action. Two of the points of 


48 Text-Book of Inorganic Chemistry . 

attraction of nitrogen, the atom of which is pentad in ammonium 
chloride, are latent in ammonia, in which this element plays the 
part of a triad. The addition of hydrochloric acid, or indeed of 
any acid, to the ammonia suffices to awaken these two latent 
points of attraction and to convert the compound of triad nitrogen 
into one of the pentad element. 

In other cases the action of heat or light is also necessary to 
make the latent points of attraction active. The four points of 
attraction of the atom of carbon, of which in the compound with 
oxygen called carbonic oxide only two are saturated, obtain their full 
value by combination with oxygen to form carbonic acid at a high 
temperature, or by combination with chlorine, producing phosgene 
gas under the influence of direct sunlight. 

We say that sulphur plays the part of a dyad element, or, more 
briefly, that it is a dyad in sulphuretted hydrogen, that it is a 
tetrad in sulphurous acid, and a hexad in sulphuric acid ; and we 
speak of dyad carbon in carbonic oxide, of triad nitrogen in 
ammonia without touching the question of what may be the 
highest atomicity of these elements. We simply pay regard to the 
results of experiment, according to which the elements unite with 
one another in certain multiple proportions. 

The chemical compounds which are produced by the union of 
the atoms are mechanically indivisible, but can be decomposed 
into their constituents by chemical forces. The smallest quan¬ 
tity of a chemical compound which cap exist in the free state is 
called a molecule. By a molecule of water we mean the compound 
of two atoms of hydrogen with one atom of oxygen ; by a molecule 
of carbonic acid, the compound of one atom of carbon with two 
atoms of oxygen. 

The molecules of water, carbonic acid, iron sulphide, &c., are 
just as little perceptible or visible as the atoms. What we call 
water is an agglomeration of water molecules, and in the same 
manner the gaseous or solid carbonic acid collected in a vessel 
consists of a large number of carbonic acid molecules. 

The question may here be asked, Are the elements as we know 
them— e.g. ordinary hydrogen gas or solid sulphur—agglomerates of 
their respective atoms ? Not long ago this was thought so; we 
now know that the atoms of the elements can combine not only 
with the atoms of other elements, but also with themselves. 

When atoms of hydrogen are set free from a chemical com- 


49 


The Atomic Theory . 

pound, and when no other substance is presented to them for 
which they possess a stronger affinity, they unite with one another 
to form molecules of hydrogen. By the chemical union of two 
atoms of hydrogen a molecule is produced, and in the same manner 
a molecule of oxygen consists of two atoms of this element, a 
molecule of nitrogen of two atoms, &C . 1 In the molecules pro¬ 
duced by the union of the atoms of the same element, the homo¬ 
geneous atoms are not bound together with less, but often indeed 
with greater, affinity than in the compounds of heterogeneous 
elements. For example, the molecule of nitrogen is a much more 
stable compound of the two atoms of the element than the com¬ 
pounds of nitrogen with oxygen, or with hydrogen or chlorine. 

Molecules of the same substance never exercise chemical ac¬ 
tion upon one another. Chemical combination or decomposition 
always occurs between the constituents of the molecules of dif¬ 
ferent bodies. Two molecules of hydrogen act upon one another 
as little as do two molecules of chlorine, but if hydrogen and 
chlorine are mixed, and exposed to the action of light or heat, a 
decomposition of the molecules of both elements ensues, with the 
mutual exchange of the atoms of each. Two molecules of hydro¬ 
chloric acid are produced from one molecule of hydrogen and one 
of chlorine. And although in this case, as in all others, it appears 
to be the molecules which act upon one another, it is in reality the 
atoms which unite with one another, and which are separated from 
molecular combination at the moment when the reaction com¬ 
mences. 

We know from experiment that the elements at the moment of 
their liberation from a compound act more energetically upon sub¬ 
stances with which they come in contact than in their ordinary 
(molecular) condition. Hydrogen and sulphurous acid in the dry 
state or in the presence of water have no action at all upon one 

1 Different molecules may exist of those elements which have a higher 
atomicity. It is possible that, besides the molecule of nitrogen, which perhaps 
consists of the two atoms of triad nitrogen united by their three points of 
attraction, a second kind of nitrogen molecule may exist in which the two 
atoms are united together by five points of attraction. Up to the present we 
only know one kind of nitrogen ; but phosphorus, sulphur, charcoal, and other 
elements exist in different forms or modifications. It is possible that the 
ordinary phosphorus, which is so easily ignited, owes its difference from the red 
amorphous phosphorus, a body endued with much less chemical affinity, to 
some such cause. 


E 


50 Text-Book of Inorganic Chemistry . 

another, but if we add sulphurous acid to a mixture of zinc and 
sulphuric acid, which evolves hydrogen, the hydrogen at the 
moment when it is set free (in the nascent state ) exercises a 
chemical action upon the sulphurous acid: it separates the mole¬ 
cules of this compound, uniting with the sulphur to form sulphu 
retted hydrogen, and with the oxygen to form water. 

We explain this and all similar phenomena by the hypothesis 
that in the nascent state the atoms, not the molecules, exert the 
chemical action, and that the atoms, with their entire force, natu¬ 
rally produce chemical changes which the elements in their mole¬ 
cular state—after the elementary atoms have united together— 
cannot effect. 


AVOGADRO’S LAW. 1 

We have here to refer to a very remarkable law concerning the 
number of molecules contained in different gaseous substances. 
This law, which is proved from the physical properties of gases 
may be thus stated :— 

Equal volumes of all gases , both elementary and compound , when 
under the same physical conditions — i.e. at the same temperature and 
under the same pressure—contain an equal number of molecules. 

It follows from this that molecular weights of all gases must 
occupy equal volumes when under the same conditions, or, in 
other words, that their molecular weights must be in the same pro¬ 
portion as their densities. If, then, we know the density of a gas 
or vapour, whether simple or compound, we can at once find its 
molecular weight, and conversely, if we know its molecular weight 
its density may be easily calculated. 

The molecular weight of hydrogen, which consists of two atoms 
and has therefore the relative weight 2, is taken as the normal 
volume, and is said to occupy two volumes. The molecular weight 
of every other gaseous substance will therefore also occupy two 
volumes under the same physical conditions, and to this law there 
are very few exceptions. The molecular weight of oxygen is 16 x 2, 
and this gas should therefore be 16 times as heavy as hydrogen, 
which agrees with experiment (sp. gr. of hydrogen = 0*0692 ; sp. gr. 
of oxygen = 1*1056 = 16 x 0*0692). Similarly, the molecular weight 

1 Introduced by the Editor. 



Avogadrds Law. 


5i 


of water, consisting of two atoms of hydrogen and one of oxygen, 

is 2x1 + 16=18, whence water gas should be — =9 times as heavy 

as hydrogen; and, finally, the molecular weight of hydrochloric 
acid, consisting of one atom of hydrogen and one of chlorine, is 

1 + 35'5=36'5) the gas should therefore be ^-5 = 18*25 times as 

2 

heavy as hydrogen. These and innumerable other results also 
agree closely with experimental determinations. 

Further, the weight of a litre of hydrogen being known 
(0*0895 gramme at o° and 760 mm.), we can easily find the weight 
of a litre of any other gas at the normal temperature and pressure, 
provided we are acquainted with its molecular weight. Thus 
the weight of a litre of oxygen under normal conditions is 

32 x 8 

0*0895 x 2 = x *43 2 gramme; of water-gas, 0*0895 x — =0806 


gramme ; of hydrochloric acid gas, 0*0895 x ='■ = 1 ‘633 gramme. 

2 

The densities of gases and vapours are usually compared with 
that of air as a standard— i.e. the density of air is taken as unity. 
In order to find from these densities the various molecular weights 
we must multiply by 14*44, which is the relative density of air com; 


pared with hydrogen ( 0 . 0 5^ 2 )> an d a £ a i n by 2, because this is the 

molecular weight of hydrogen. In other words, we multiply by 
28*88. The density of water gas, for example, is 0*622, and its 
molecular weight 0*622 x 28*88 = 17*96, or very nearly the same as 
its theoretical molecular weight (18). 1 


CHEMICAL AFFINITY AND UPON WHAT IT 
DEPENDS. 

The chemical force of attraction (affinity) which resides in the 
atoms of the elements is of different strength. For example, the 
affinity of potassium for chlorine is so much greater than it is for 
iodine, that if chlorine is led into a solution of potassium iodide 

1 For a more detailed account of Avogadro’s law, the student is referred to 
the Appendix.—E d. 


E 2 




52 Text-Book of Inorganic Chemistry. 

’the chlorine unites with the potassium, and all the iodine is set 
free. In the same manner, the affinity of iron for chlorine is so 
■much greater than that of copper, that if a bright rod of iron is 
placed in a solution of copper chloride, it becomes at once covered 
with a coating of metallic copper, and at last all the copper is 
separated out from the compound in the metallic form, while an 
equivalent quantity of iron unites with the chlorine which was pre¬ 
viously combined with the copper. 

But, it may be asked, How much greater is the affinity of 
potassium for chlorine than for iodine, or of chlorine for iron than 
for copper ? And it might be thought that the experimental answers 
’to such questions ought not to be more difficult than the measure¬ 
ment of the force of gravity, the strength of an electric current, and 
-other physical forces. But chemistry is in this respect far behind 
physics, and not only because chemistry is a younger science, but 
also because chemical affinity is largely influenced by a number of 
’important conditions which often occur together—by temperature, 
light, electricity, state of aggregation, quantity, &c., and which 
often cause a powerful affinity to be weakened, and a weak affinity 
to be strengthened. Some examples of this may serve to show how 
difficult it is to determine the force of these oft-changing chemical 
affinities. 


Influence of Temperature on Chemical Affinity. 

At the ordinary temperature, mercury and oxygen do not 
■possess sufficient affinity to combine with one another, nor is their 
affinity strong enough to produce this change at the boiling-point 
of water ; but at a temperature of about 300° they unite chemically 
with one- another, and form red oxide of mercury. This compound 
is, however, again decomposed into its constituents at a slightly 
higher temperature (under 400°). At 400° mercury and oxygen have 
as little affinity for one another as at the ordinary temperature. 
The metal potassium, when moderately heated in a stream of car¬ 
bonic acid gas., abstracts the oxygen from this 'Compound to form 
an oxide, and then a carbonate, charcoal being separated. But if 
potassium carbonate is heated with charcoal up to bright redness, 
the affinity of the charcoal for the oxygen becomes greater than 
that of the metal for the same element, and the potassium is again 
set free. Charcoal, which at the ordinary temperature has a very 
weak affinity for oxygep, obtains the most powerful .affinity for this 
element at a red heat. 


Chemical Affinity , and upon what it depends . 55 


Influence of Light on Chemical Affinity. 

A mixture of about equal volumes of hydrogen and chlorine, 
when quite excluded from the light, remains unchanged, and from' 
this we conclude that the atoms of these elements in their molecules 
do not possess sufficient affinity to combine with one another. But 
if the mixture is exposed to light, the combination of the two gases, 
producing hydrochloric acid, goes on the more quickly the more 
intense the light. It is, further, principally the violet rays of light 
which produce this chemical change ; and these rays are therefore 
called the chemically active ( actinic ) rays. If the mixture contains 
pure hydrogen and pure chlorine in exactly equal volumes, diffused 
daylight produces instantaneous combination. Light is thus able 
to awaken the affinity of hydrogen and chlorine for one another, 
and so to cause them to combine. On the other hand, stable com¬ 
pounds— e.g. those of silver with chlorine or bromine—are some¬ 
times decomposed by light. We thus see that light also can cause 
both chemical combination and decomposition. 

Influence of the State of Aggregation and of Quantity (Mass) 
on Chemical Affinity. 

Potash has a stronger affinity for sulphuric acid than lime, and 
sulphuric acid has a stronger affinity for bases than acetic acid ; 
but notwithstanding this, if we mix a solution of potassium sulphate 
with one of calcium acetate, calcium sulphate separates out and 
potassium acetate remains in solution. In this case it is the insolu¬ 
bility of calcium sulphate (which is nearly insoluble in water) that 
causes the stronger acid (sulphuric acid) to separate itself from the 
stronger base—potash—and to unite with the lime. 

In the same manner, the quantities or masses of different bodies 
acting upon one another also influence chemical changes. Sulphuric 
acid has a much stronger affinity for metallic oxides than hydro¬ 
chloric acid, still the latter can partially displace the former in 
solutions of the sulphates, and the more completely the more it is 
in excess. If, for example, strong hydrochloric acid in considerable 
quantity is added to the blue aqueous solution of copper sulphate, 
free sulphuric acid is obtained with copper chloride, the latter 
being easily recognized by the green colour which the blue solution 
gradually acquires. 

These examples suffice to show that the force of chemica 


54 Text-Book of Inorganic Chemistry. 

affinity residing in the atoms of the element is difficult to measure, 
since it is influenced and changed in very various ways by other 
secondary forces and conditions. Before we are in a position to 
measure accurately the strength of these secondary forces, even 
when several act together, we shall not be in a position to predict 
accurately and to calculate mathematically what would be the re¬ 
sult when substances, acting chemically upon one another, are 
brought into contact under different conditions ; although this is 
always the ultimate goal of all chemical investigation. 

But as long as we remain far distant from this end, as at present, 
chemistry is a science in which mathematical calculations can be 
little used. 1 


CHEMICAL NOMENCLATURE. 

The impossibility of inventing empirical names for many 
thousand chemical compounds has been long admitted, and at 
an early date it was proposed to form the names so that everyone 
might easily understand the nature of any particular compound, 
and to what class it belonged, by the name given to it. 

But besides these rational designations, other empirical names, 
which come to us with the authority of age, and which also possess 
the merit of being shorter, are still employed, especially by those 
unacquainted with the principles of chemistry. The chemist 
occasionally makes use of the names : soda, potash, Glauber’s salts, 
lunar caustic, &c., although in inorganic chemistry such names are 
but little employed. In organic chemistry, on the other hand, 
where we have to refer to whole classes of bodies, of whose chemical 
composition we know but little, as sugar, starch, the glucosides, 
&c., we must content ourselves for the present with the empirical 
names. 

1 Other things being equal, the strength of the affinity which causes two 
substances to unite with one another is proportional to the quantity of heat 
set free on their combination. The quantity of heat evolved when a molecule 
of any compound is produced under the same physical conditions is always the 
same, and if certain changes are possible among a system of bodies, that 
change will take place which produces the greatest amount of heat. Thus, if 
two elements, A and B, are mixed with a third, C, with which each can unite 
directly, combination will take place between those two which liberate the 
greater quantity of heat on their union.—E d. 



Chemical Nomenclature . 


55 


The compounds produced by the union of the elements, and of 
these compounds again combined together, mostly belong to three 
classes— bases , acids , and salts. The compounds of oxygen with the 
other elements, the most numerous of all, are called oxides ; in the 
same manner, those of sulphur are sulphides , those of chlorine are 
chlorides , of bromine bromides , &c. 

The oxides may be divided into two great classes —basic oxides 
and acid oxides ; those not belonging to either of these classes are 
mostly indifferent compounds. 

Those oxides which have a basic character— i.e. which combine 
with acids to form salts—may be called the oxides proper, and 
their name is built up from the name of the element with which 
the oxygen is combined. Thus we say potassium oxide, zinc 
oxide, &c. 

In some cases we are acquainted with two basic oxides of an 
element. We know, for example, two basic oxides of copper, iron, 
mercury, and other metals. In order to clearly distinguish these 
from one another, the termination of the metal is changed : the 
termination -ic is added to the metal to indicate the compound con¬ 
taining the greater quantity of oxygen, and the termination -ous to 
indicate the compound containing the smaller quantity of oxygen. 
Thus we say mercuric oxide and mercurous oxide. In cases where 
these terminations when added to the name of the metal would be 
harsh, the Latin name of the metal is employed : we say, for 
example, ferric oxide and ferrous oxide for the two oxides of iron, 
cupric oxide and cuprous oxide for those of copper, &c. 

Those basic oxides, as ferric oxide, chromic oxide, which contain 
the atoms of the metal, and oxygen in the proportion of two to three, 
are often called sesqui-oxides , or one-and-a-half oxides. 

Oxides of the metals, which possess neither acid nor basic 
character, and which contain more oxygen than the basic oxide, 
are called peroxides. We are acquainted, for example, with man¬ 
ganese peroxide, lead peroxide, &c. They give up the excess 
of oxygen which they contain when heated with strong acids. 

Suboxides, of which the number is very small, are those indif¬ 
ferent oxides of the metals which contain less oxygen than the 
basic oxides. On treatment with acids they yield the metal and 
the basic oxide, which is richer in oxygen. Such an oxide is lead 

The oxygen compounds which are said to have an acid charac¬ 
ter, and are sometimes called acids, are those which unite with 


5 6 Text-Book of Inorganic Chemistry. 

bases to form the ordinary or oxygen salts. They form with water 
the true oxy-acids, and are called atihyclrides. The names are built 
up in the same way as those of the oxides, except that the termi¬ 
nation -ic is always added if there be only one compound of the 
element with oxygen. Thus, the single oxide of carbon of an acid 
character is called carbonic anhydride. 

We further distinguish between those anhydrides of the same 
element which contain different quantities of oxygen by the same 
terminations as are used for the oxides (-ic and - ous ) ; for example, 
the two anhydrides of sulphur are sulphuric anhydride and sul¬ 
phurous anhydride; of nitrogen, nitric anhydride and nitrous 
anhydride. 

Some elements—^, chlorine—are capable of forming a larger 
number of oxides with acid characteristics. These are distinguished 
by the prefixes per- and hypo-. The oxide which contains more 
oxygen than chloric anhydride is called perchloric anhydride, that 
which contains less than chlorous anhydride is called hypochlorous 
anhydride. 

The oxy-salts are named from the bases and acids of which 
they consist. Salts formed from acids with the termination - ic, 
have the termination -ate, while those from acids with the termina¬ 
tion -ous end in -ite. For example, potassium sulphate is the com¬ 
pound of potassium oxide and sulphuric acid ; cupric nitrate, of 
cupric oxide and nitric acid ; mercurous nitrite, of mercurous oxide 
and nitrous acid. 

The name hydrate includes two classes of bodies : the acid 
hydrates (the oxy- or sulpho-acids) and the basic hydrates (the 
hydrates proper). The acid hydrates, or as they are better called 
acids, and which are produced by the union of the anhydrides 
with water, usually contain in their molecule the same number of 
hydrogen atoms as the normal salt of a monad metal contains of a 
metal. 1 The names of the acids are derived from those of the 
corresponding anhydrides ; thus, from sulphuric anhydride we get 
sulphuric acid, from phosphoric anhydride, phosphoric acid, &c 

Basic hydrates are the oxides of the metals in which a portion 
of the metal is displaced by hydrogen— e.g. potassium hydrate, 
cupric hydrate. They are often produced by the union of the 

1 An acid is therefore a hydrogen salt, and the hydrogen which can be 
displaced by a metal, not necessarily all which the acid contains, is called its 
displaceable hydrogen. Accordingly, as an acid contains one, two, three, &c., 
atoms of displaceable-hydrogen, it is said to be mono-, di-. tribasic, &c.—E d. 


Chemical Nomenclature. 5 7 

basic oxides with water, and may be considered, as was formerly 
the case, as compounds of these two substances. 

The chemical nomenclature of the acids, bases, and salts, which 
contain sulphur instead of oxygen is not so well developed, and 
probably because we are less accurately acquainted with these 
compounds. They may be called sulpho-acids , sulpho-bases , and 
sulpho-salts. On the union of a'sulpho-acid ( e.g . carbon disul¬ 
phide or arsenious sulphide) with a sulpho-base {e.g. potassium 
sulphide or sodium sulphide) a sulpho-salt is produced. Thus 
carbon disulphide (sulpho-carbonic acid) with sodium sulphide 
gives sodium sulpho-carbonate; arsenious sulphide (sulph-arsenious 
acid) with potassium sulphide gives potassium sulph-arsenite, and 
in this way all other sulpho-compounds are distinguished. 

The salts which result from the direct union of the halogen 
elements (chlorine, bromine, iodine, and fluorine) with the metals, 
and which contain neither oxygen nor sulphur, are called the 
haloid-salts , and the corresponding acids, containing hydrogen in 
the place of the metal, are the haloid-acids. 

•Hydrochloric acid, the compound of hydrogen and chlorine, 
is not to be confused with chloric acid, which contains oxygen as 
well as these elements. 

The metals unite with the halogens usually in the same pro¬ 
portion as with oxygei>, and the corresponding names are the 
same. To mercuric oxide corresponds mercuric chloride, con¬ 
taining a quantity of chlorine equivalent to the oxygen in the 
oxide ; similarly, mercurous chloride corresponds to mercurous 
oxide. 

In cases when elements unite with the halogens in more than 
two proportions, or when, as in the case of phosphorus, the com¬ 
pounds cannot be included under the true haloid-salts, it is 
customary to indicate the number of chlorine, bromine, &c., atoms 
contained in the molecule of the compound by the prefixes 
mono-, di -, tri-, ietra -, penta -, hexa -, &c. The two compounds of 
phosphorus and chlorine, which contain respectively five and three 
atoms of chlorine united to one atom of phosphorus, are called 
phosphorus pentachloride and phosphorus trichloride; and, in 
the same manner, we also say carbon disulphide, referring to the 
compound containing one atom of carbon and two of sulphur in 
the molecule. 

Haloid compounds are also known, in which a portion of the 
chlorine, bromine, &c., is displaced by oxygen. These are called 


58 Text-Book of Inorganic Chemistry. 

oxychlorides , oxybromides , &c. To this class belongs the well- 
known phosphorus oxychloride. Those compounds in which 
sulphur displaces a portion of the halogen instead of oxygen are 
similarly designated—^, the compound of phosphorus, chlorine 
and sulphur is called phosphorus sulphochloride. 

Alloys are the intimate mechanical mixtures produced by 
melting metals together. Those alloys which contain mercury are 
called amalgams. Alloys and amalgams can only be produced 
from metals. 


CHEMICAL SYMBOLS AND FORMULAE. 

When we calculate with figures we do not use the words 
representing them, but their signs. In the same manner, the 
need of symbols to express chemical substances, and by a suitable 
combination of these symbols to reproduce chemical thoughts in a 
brief but general manner, was early recognized in chemistry. 

It is customary to represent the elements by letters and to 
express chemical compounds by the juxtaposition of these letters. 
The symbol H has been chosen to represent the element hydrogen, 
and I the element iodine ; by placing these symbols together, 
without any sign between them, the compound of the two elements 
—hydriodic acid—is represented (HI). In the same way the 
compound symbol CaO is a simple expression for the compound 
of the metal calcium (Ca) with the gas oxygen (O )—i.e. common 
quick-lime. 

The symbols of the elements are the first letter of their names 
and usually of their Latin names : thus K represents potassium 
(kalium), C = carbon, &c. When the names of several elements 
begin with the same letter, some other letter is added : thus B 
means boron, Ba = barium, Br - bromine, Mg = magnesium, Cd = 
cadmium, Ag = silver (argentum), Hg = mercury (hydrargyrum), 
Fe = iron (ferrum), Sn = tin (stannum), &c. 

But the chemical symbols of the elements have another and 
more important meaning. The symbol HI does not only indicate 
the presence of hydrogen and iodine, but expresses also the 
quantities by weight of the constituents of hydriodic acid which are 
contained in its molecule. The symbol H indicates one atom of 
hydrogen, or the smallest weight which enters into a chemical 



Chemical Symbols and Formulce. 59 

compound (which is taken as unity), and, in the same manner, I 
means one atom of iodine or 127 parts by weight; while, finally, the 
compound symbol HI expresses the compound of hydrogen and 
iodine, which contains one part of the former united with 127 
parts of the latter, and of which the molecular weight is, therefore, 
127 + 1 = 128. 

The accompanying table on p. 60 gives the names of the 
elements, alphabetically arranged, with their symbols and atomic 
weights, the non-metallic elements being in italics. 

Since, however, the elements combine with one another not 
only in simple but also in multiple proportions, and as this must 
also be represented by chemical symbols, it is customary to ex¬ 
press the number of atoms of the one element which unites with 
the other to form a molecule by a small figure placed to the right 
of the symbol of the element. For example, S0 2 means the 
compound of one atom (32 parts) of sulphur with two atoms 
(2x16 parts) of oxygen, and the formula H 2 0 means the compound 
produced by the union of two atoms of hydrogen (2 parts) with one 
atom of oxygen (16 parts). 

The figures placed to the right of any chemical symbol refer 
only to this one symbol, but another large figure placed on the left 
of a symbol multiplies all that follows up ^o a comma, full-stop, or 
plus sign. The formula 2KCl,PtCl 4 , for example, represents the 
compound of two atoms of potassium, and two atoms of chlorine 
(or two molecules of potassium chloride), with one of platinic 

chloride. Similarly, the formula S0 2 j + i°H 2 0 is the 

symbolic expression for a compound of sodium sulphate with 10 
molecules of water— i.e. with 20 atoms of hydrogen and 10 atoms 
of oxygen. 

Every such formula is full of hidden meaning to the chemist. 
The cbmposition of the above compound (Glauber’s salts) might 
be more briefly expressed by the formula SNa 2 H 20 O 14 , but this 
only tells us that one molecule contains one atom of sulphur, two 
of sodium, twenty of hydrogen, and fourteen of oxygen. If we 
wished to express the fact that ten molecules of water are 
contained in each molecule of the salt, we might write its formula 

S 0 4 Na 2 + ioH 2 0. But the above formula, S 0 2 j q N ^ + ioH 2 0, 

in which the constituents of the compound are still further arranged, 
expresses the fact that the four oxygen atoms in the sodium sub 


6 o 


Text-Book of Inorganic Chemistry. 


Table of Atomic Weights. 


Name 

O 

6 

>-* 

in 

Atomic 

weight 

Aluminium 

A 1 

27 

Antimony . 

Sb 

120 

A rsenic 

As 

75 

Barium 

Ba 

137 

Beryllium . 

Be 

13-6 

Bismuth . 

Bi 

210 

Boron . . j 

B 

11 

Bromine . . | 

Br 

80 

Cadmium . 

Cd 

112 

Caesium . 

Cs 

133 

Calcium . 

Ca 

40 

Carbon 

C 

12 

Cerium 

Ce 

141 

Chlorine . 

Cl 

35-5 

Chromium 

Cr 

52-2 

Cobalt 

Co 

59 

Copper 

Cu 

63-4 

Didymium 

Di 

146 

Erbium 

Er 

166 

Fluorine . 

F 

19 

Gallium 

Ga 

69-8 

Gold . 

Au 

197 

Hydrogen . 

H 

1 

Indium 

In 

ii 34 

Iodine 

I 

127 

Iridium 

Ir 

193 

Iron . 

Fe 

56 

Lanthanum 

La 

139 

Lead . 

Pb 

207 

Lithium 

Li 

7 

Magnesium 

Mg 

24 

Manganese 

Mn 

55 

Mercury . 

Hg 

200 


Name 

"3 

g 

to 

Atomic 

weight 

Molybdenum 

1 Mo 

! 

96 

Nickel . . | 

Ni 

58-5 

Niobiujn . . 1 

Nb 

94 

Nitrogen . 

N 

14 

Osmium . 

Os 

199 

Oxygen 

0 

16 

Palladium. 

Pd 

106 

Phosphorus 

P 

3 i 

Platinum . . | 

Pt 

195 

Potassium. 

K 

39 

Rhodium . 

Rh 

104 

Rubidium . 

Rb 

85*4 

Ruthenium 

Ru 

104 

Selenium . 

Se 

79 

Silicon . . ; 

Si 

28 

Silver . . j 

A g 

108 

Sodium . . j 

Na 

23 

Strontium . . ; 

Sr 

87-5 

Sulphur . . I 


32 

Tantalum . 

Ta 1 

182 

Tellui ium . . | 

Te 

128 

Terbium . 

Tr 

148 

Thallium . 

T 1 

204’ 

Thorium . 

Th 

232-5 

Tin . 

1 Sn 

i 118 

Titanium . 

Ti 

48 

Tungsten . 

W 

184 

Uranium . 

\ u 

240 

Vanadium 

V 

5i-3 

Yttrium 

Y 

90 

Ytterbium. 

Yb 

173 

Zinc . 

Zn 

65 

Zirconium . 

Zr 

1 

90 






































Chemical Symbols and Formnlce. 6 1 

phate perform different functions. Two of them are considered to 
be in closer connection with the atom of sulphur than the other 
two. These two atoms of oxygen saturate four of the bonds of the 
hexad sulphur atom, producing a dyad group of elements (SO a ) 
which is capable of playing the part of a single element, and 
which is called a compound radical. There is another fact ex¬ 
pressed in this formula—viz. that this dyad radical does not unite 
directly with the two atoms of sodium, but through the intervention 
of the other two atoms of oxygen, which have, therefore, other 
functions than the two former atoms of this element. 

The employment of chemical symbols makes the meaning of 
chemical decompositions much more easy to understand. And the 
use of formulae enables us to determine what quantity of a body is 
necessary to produce, either alone or by its action on a second 
substance, a given quantity of a third. 

We have seen (p. 12) that potassium chlorate breaks up on 
heating into potassium chloride and oxygen. In order to find out 
how much of these two substances would be obtained from 100 
grammes of potassium chlorate, we must first know the molecular 
weight of this salt, and how many atoms of each of its elements 
are contained in its molecule. We could, of course, express this 
in words, and say one molecule of potassium chlorate consists of 
one atom of potassium, one of chlorine and three of oxygen, and, 
the atomic weights of these elements being known, that one mole¬ 
cule of the salt contains :— 

35-5 x 1 = 35*5 parts by weight of chlorine, 

39-0 x 1 = 39*o „ potassium, 

1.6 xj = 48-0 „ » oxygen, 

122-5 » » potassium chlorate, 

which means that on the decomposition of 122-5 parts by weight 
of potassium chlorate 48 parts of oxygen would be given off, and 
74 - 5„parts of potassium chloride would remain behind. 

But all this, which is expressed by so many words, may be at 
once represented by the simple chemical equation 

CIO 3 K = KC1 + 3O. 

Cl means 35*5 parts by weight of chlorine, 3O means 3 x 16 = 48 
parts of oxygen, and K 39 parts of potassium. Placing these 
numbers opposite the symbols we get:— 



62 


Text-Book of Inorganic Chemistry. 


C 10 3 K = KC 1 + 3O, 

35-5+48 + 39 = (39 + 35*5) + 48, 

122-5 = 74*5 + 48, 

which shows that 122-5 parts by weight (grammes, &c.) of potas¬ 
sium chlorate yield 74-5 parts (grammes, &c.) of potassium chloride, 
and 48 parts (grammes, &c.) of oxygen. From 100 grammes of potas¬ 
sium chlorate we should, therefore, obtain the weights of potassium 
chloride (x) and oxygen (y) expressed in the two following simple 
rule-of-three sums :— 


122*5 : 74*5 :: 100 : x x = 6 o -8 grammes potassium chloride. 

122-5 : 48 :: 100 : y .*.JK = 39'2 grammes oxygen. 

This simple method gives a ready means of calculating the 
weight of any particular substance which can be obtained from a 
given weight of some other substance in a known chemical reac¬ 
tion. Suppose, however, we wish to know the volume of the sub¬ 
stance produced when it is a gas. Let us take the same example 
and imagine that it is required to find how many litres of oxygen 
would be produced when 100 grammes of potassium chlorate were 
completely decomposed. We have seen above that the weight of 
oxygen produced is 39-2 grammes, and we further know that the 
weight of a litre of oxygen is 16 times that of a litre of hydrogen 
(the molecule of the former being 16 times as heavy as that of the 
latter), or, in other words, is 0*0895 x 16= 1-432 grammes. From 
this, therefore, it follows that the 39-2 grammes of oxygen will 

occupy a volume of ^ 9 2 =27-3 litres. This is of course the 
** 1 * 43 2 

volume of the gas at the normal temperature (o° C.) and under 
the normal pressure (760 mm.). In the same manner the volume 
of any given weight of a gas (and the converse) may be easily 
found, provided we know its molecular weight. 

It has been previously stated (p. 9) that manganese peroxide 
when heated gives off oxygen. The process, which has been 
shown by analytical experiments to be according to the following 
equation, is somewhat less simple. From this equation— 

3Mn0 2 = Mn 3 0 4 + 0 2 , 


we learn that three molecules (261 parts by weight) of the per¬ 
oxide yield one molecule (229 parts) of trimanganic tetroxide, and 
one molecule (32 parts) of oxygen. Whence it follows that 12*2 
kilos, of oxygen will be obtained from 100 kilos, of the peroxide. 

The chemical manufacturer who wishes to convert 100 kilos, of 



General Remarks on the Chemical Elements. 63 

saltpetre (potassium nitrate) into nitric acid, so as to obtain the 
largest yield of acid possible, and to use the smallest quantity of 
sulphuric acid necessary, must of course know the details of the 
process he is about to employ, and must remember that acid 
potassium sulphate is obtained as a bye-product. The reaction 
which takes place is expressed by the following equation :— 

NOj • OK + S 0 2 |Og = SO.jgj* + NOj • OH. 

Potassium Sulphuric Acid potassium Nitric 

nitrate acid sulphate acid 

Translating the formulas into numbers we get— 


N = 14 

S = 32 

S 

= 32 

N 

= 14 

0 3 = 48 

o 4 = 64 

O, 

= 64 

O3 

= 48 

K = 39 

H 2 = 2 

K 

« 39 

H 

= 1 



H 

= 1 



IOI 

98 


136 


63 

which show that 

the manufacturer 

must 

use 98 kilos, of sulphuric 


acid for every 101 kilos, of saltpetre, and that he will obtain 136 
kilos, of acid potassium sulphate and 63 kilos, of nitric acid. 


GENERAL REMARKS ON THE CHEMICAL 
ELEMENTS. 

It is remarkable that the elementary constituents of the earth, 
of which we know more than sixty, occur both in the free state 
and combined in very different quantities and are very differently 
distributed on the earth. Some may be found everywhere, others 
occur rarely and in small quantities ; while others, again, are found 
only in a few places, but then in considerable quantities. The most 
widely distributed elements are oxygen, silicon, aluminium, iron, 
and calcium ; those which are more scarce include the noble 
metals gold, platinum, &c. Others, which only occur in definite 
places, but then in considerable quantities, are such as mercury, 
tin, &c. Finally, some are so rare that we only possess a very 
meagre knowledge of them and their compounds— eg. indium, 
caesium. 

Some elements occur as such free in nature, or, as it is said, 
native — eg. platinum—but the majority of the elements are found 



64 Text-Book of Inorganic Chemistry. 

exclusively in combination with others; while some, as sulphur, 
oxygen, nitrogen, silver, copper, occur both free and combined. 

A question which has often been asked but never answered, 
and perhaps never will be, is whether the interior of the earth may 
not contain other elements besides those with which we are ac¬ 
quainted, and whether the elements and their compounds may not 
be differently distributed there than they are in the portion of the 
earth’s crust accessible to us. 

The portion of the earth which we can investigate— i.e. the atmo¬ 
sphere, the ocean, and a very thin layer of the solid crust—is only 
a minute fraction of the whole mass. The mean radius of the 
earth is over 3,950 miles, and as we descend towards the centre 
the temperature increases to such an extent that at a depth of about 
fifty miles nearly all the substances we know on the earth would 
be in a liquid state. That this is the case in some places is proved 
by the existence of volcanoes which at times pour forth streams of 
liquid lava. But it must not be forgotten that the substances 
forming the interior of the earth are under a much greater pressure 
than those at the surface, and we cannot tell how far the effect of 
this pressure would modify the melting-points of the different 
bodies. It may be considered as probable that the interior of the 
earth consists of substances heavier than those at the surface, for 
while the mean specific gravity of the rocks at the surface is be¬ 
tween 2-5 and 3*0, that of the entire earth is about 5-5. But here 
again pressure would tend to compress the substances in the 
interior and so to make them heavier ; for example, at the ordinary 
temperature, water would become as heavy as mercury at a depth 
of 360 miles. How far these two opposing forces, that of the inter¬ 
nal heat and that of pressure, counterbalance one another remains, 
and probably will remain, an unsolved problem for us. 

But it is certainly to be expected that many unknown elements 
are hidden in the interior of the earth, perhaps some even heavier 
than platinum and iridium, which, owing to their high specific 
gravity, can never approach the earth’s surface ; while, on the other 
hand, it may be mentioned that numerous analyses of meteorites 
which have fallen to the earth from extra-terrestrial space have 
not revealed any new elements to us. 

We consider the present chemical elements to be simple bodies 
because we are not able at present to further decompose them 


General Remarks on the Chemical Elements. 65 

(p. 5), but we do not therefore think them incapable of decompo¬ 
sition. On the contrary, the suspicion that many of them are 
either compound bodies or modifications of one element is sup¬ 
ported by weighty reasons. 

In the first place, the history of chemistry teaches us that sub¬ 
stances long considered as elements can be broken up by improve¬ 
ments in the methods and means of decomposition into simpler 
bodies, and now it is thought possible that even these may be 
compounds. 

Potassium hydrate, a compound of potassium, oxygen, and 
hydrogen, quick-lime, consisting of calcium and oxygen, and other 
bodies, were thought to be elements at the beginning of the present 
century, when the electric current as a powerful means for effecting 
decomposition was still unknown. By electrolysis the unknown 
metal potassium was recognized as a constituent of potassium 
hydrate, and the metal calcium was isolated from quick-lime. Such 
events justify the expectation that with the discovery of new and 
more powerful means of decomposition, many of the elements would 
disappear from the present list. 

Compound bodies are known, the compounds of which have 
so close an analogy to the similar compounds of certain elements 
that it is often difficult to separate them from one another. Such 
compounds would certainly, have been classed with the elements had 
not their decomposition been accomplished. Of these, cyanogen , 
a compound of carbon and nitrogen, is a striking example. The 
compounds of this substance with the metals and with hydrogen 
closely resemble the similar compounds of the elements chlorine, 
bromine, and iodine. And this similarity makes it quite possible 
that chlorine, bromine, and iodine will, some time or other, prove 
to be compounds. 

Besides chlorine, bromine, and iodine, other elements exist 
which possess remarkable similarity in their compounds. Such 
are, for example, the groups : sulphur, selenium, and tellurium ; 
potassium, sodium, and lithium ; barium, strontium, and calcium. 
It is certain that these similarities are just as little accidental 
as the fact that the similar elements generally occur associated 
together in nature. 

Still more remarkable are the relations in which the atomic 
weights of the elements of each group stand to one another. The 
atomic weight of the middle element is almost exactly the mean 
of those of the other two. For example, in the group potassium, 

F 


66 Text-Book of Inorganic Chemistry. 

sodium, and lithium, with the atomic weights 39, 23, 7 respectively, 
the mean of 39 and 7, or , is exactly 2 3—t.e. the atomic 

weight of sodium. Similarly in the second group, barium (137)? 
strontium (87-5), and calcium (40), the mean of 137 and 40 is 
88*5, or very nearly the atomic weight of strontium ; in the same 
manner in the third, chlorine (35*5), bromine (80), and iodine 
(127), the mean of 35-5 and 127 (81*25) nearly coincides with the 
atomic weight of bromine (80); and finally, in the fourth group : 
sulphur (32), selenium (79), and tellurium (128), the mean of the 
atomic weights of sulphur and tellurium is 80, or very nearly the 
same as the atomic weight of selenium (79). It will be noticed 
that the numbers do not exactly agree, although they very nearly 
do so. 

It is certainly not accidental that the members of these natural 
groups and other similar ones always occur together in nature, 
that they have a close chemical likeness to one another, and that 
such a remarkable relation connects their atomic weights together. 
We do not know the cause and connexion of these facts ; we can 
only guess and surmise that truths lie hidden in them which, if 
brought to light, would considerably enrich and further our 
science. 

The investigator cannot attempt to explain every fact, even 
those close at hand, and must beware of substituting speculation 
for experiment. Much might be written on facts such as those to 
which we have referred, and brilliant hypotheses might be built 
upon them ; but these theories often lose their splendour when ex¬ 
perimental investigation lifts the veil and discloses the truth. 

As might be expected from the great divergence in the atomic 
weights, similar groups can be easily built up from other elements 
than those mentioned above. The atomic weights of oxygen, 
nitrogen, and carbon offer an example. The mean of the atomic 
weights of oxygen (16) and carbon (12) is 14, or exactly the atomic 
weight of nitrogen. Still no one would think of placing these 
elements together in a group because of the great difference in 
their chemical behaviour. The compounds of oxygen, nitrogen, 
and carbon with other elements— e.g. with hydrogen—have not 
the slightest resemblance to one another; they may in fact be said 
to differ from one another as much as the hydrogen compounds of 
the natural group chlorine, bromine, and iodine resemble one 
another. These three elements cannot, therefore, be grouped 



Physical and Chemical Properties op Bodies . 67 

together in the same manner as chlorine, bromine, and iodine, 
&c. : the connexion between their atomic weights appears to be 
purely accidental. 

Remarkable relations between the elements are also found in 
another direction. There are certain elements of which the 
atomic weights are almost exactly the same, and which always 
occur associated together in nature. To these belong the metals 
platinum and iridium, with nearly the same atomic weight (195 
and 193) ; rhodium and ruthenium (104); cobalt and nickel, with 
the atomic weights 59 and 58-5 ; and iron and manganese (56 and 
55), as well as some others. 

These facts obtain increased importance since those elements, 
which occur together in nature and possess about the same 
atomic weight, always exhibit a close relation in their chemical 
properties. And since we know that elements exist which, though 
chemically identical, can possess very different physical and 
even chemical properties (e.g. graphite and diamond, both modifi¬ 
cations of the element carbon ; ordinary and red or amorphous 
phosphorus), the supposition is not unsupported that at a later 
time it may be possible to prove the identity of cobalt and nickel 
or of iron and manganese, and even to convert the one into the 
other. 

From whichever side the question of the elementary nature of 
the elements is attacked, it seems probable, although no decided 
answer can be obtained, that the elements, or the greater part of 
them, are capable of decomposition and that their number might 
be much reduced. 


PHYSICAL AND CHEMICAL PROPERTIES OF 
BODIES. 

When two bodies are allowed to act upon one another it is 
easy in most cases to decide whether a chemical change has taken 
place or not. Sometimes, however, especially in organic chemistry, 
careful observations and experiments are necessary to decide 
whether two products of different reactions are identical or dif¬ 
ferent. This is done by an exact determination of the physical and 
chemical properties of the bodies under investigation, the enu¬ 
meration of which constitutes their real description. 



68 Text-Book of Inorganic Chemistry. 

One of the most important of the physical properties of a body 
is its 

State of Aggregation .—It must be known Whether the body in 
question is a solid, a liquid, or a gas under normal conditions, and 
whether it changes its state of aggregation, and if so, under what 
conditions. If the body is a liquid, its boiling-point, reduced to 
the normal atmospheric pressure, must next be determined, as 
well as the temperature at which the liquid becomes solid, if at all, 
or its freezing-point (melting-point). If the substance is a gas, the 
temperature, or pressure, or both, under which it becomes liquid 
must be found. With solid bodies the determination of the tem¬ 
perature at which it becomes liquid (melting-point) must be made. 

The colour of the substance must also be noticed; it often 
changes with the state of aggregation. Sulphur, for example, in 
the solid state is bright yellow, in the liquid state yellow to brown, 
in the gaseous state reddish-brown; iodine, at the ordinary tem¬ 
perature, is a dark-grey solid with metallic lustre, but when melted 
it becomes reddish-brown, and in the gaseous state has a splendid 
violet colour. 

The taste and odour .—The determination of these properties, 
especially the former, require caution, owing to the poisonous nature 
of many substances, and because of the caustic action of others 

_ eg. sulphuric acid, caustic potash—on the mucous membrane 

of the tongue and mouth. A large number of volatile bodies may 
be easily and tolerably certainly recognized by their odour alone— 
e.g. chlorine, iodine, sulphurous acid, chloroform, carbon disulphide, 
camphor. We distinguish different tastes as : acid (sour), alkaline 
(soapy), acrid, astringent, sweet, bitter, insipid, &c.; and different 
odours as : acid, putrid, suffocating, &c. The odour of bodies is 
often compared with that of other well-known substances ; thus, 
we say an ethereal odour, or that a body smells like cinnamon, 
bitter almonds, camphor, butyric acid, &c. 

Lustre , Fracture , and Hardness .—Among solids, different kinds 
of lustre of their surfaces are distinguished according to substances 
whose lustre is well known— e.g. metallic lustre, diamond lustre, 
vitreous lustre, silky, mother-of-pearl lustre, &c. 

The fracture of a solid when broken is indicated by the same 
terms as the lustre. Some bodies have a crystalline fracture, others 
a conchoidal (shell-like), or a splintery fracture. The crystalline 
fracture may be laminated, granular, or fibrous. 


Physical and Chemical Properties of Bodies. 69 

For the determination of the hardness of solids, the same series 
of minerals is used as in mineralogy. This scale consists of a 
series of ten minerals of different degrees of hardness, commencing 
with the diamond and ending with talc. Liquids are distinguished as 
viscid and mobile. It often happens that viscid liquids are heavy 
— e.g. concentrated sulphuric acid—and that mobile liquids are 
light, although the greater or less consistency of the liquid does 
not depend upon its specific gravity. The heavy liquid chloroform 
is very mobile, while the light fatty oils are viscid. 

The specific gravity or deitsity of a substance indicates how 
many times heavier the substance is than some other known body. 
Water is chosen as the unit for liquids and solids, air as the unit 
for gases and vapours. 

A piece of iron weighed in the air is much heavier than when 
dipped in water. In the latter case its loss in weight is exactly 
the weight of the water which it displaces. The absolute weight of 
the iron divided by the weight of the water displaced (its own 
volume) gives the specific gravity of iron (7*8). Similarly deter¬ 
mined, the specific gravity of platinum is nearly three times as 
much (21 *5). 

If we take a small vessel, of known weight, with a narrow neck 
on which is a mark, fill it with mercury and weigh it, and then do 
the same with water, we obtain, on subtracting the weight of the 
flask from each weighing, the relative weights of mercury and 
water (equal volumes), and, dividing the former by the latter, the 
specific gravity of mercury (13-6) is found. 

In a similar manner the specific gravities of gases and vapours 
may be obtained, employing the word vapour for the gases pro¬ 
duced on heating volatile liquids above their boiling-points. 

The methods by which determinations of specific gravity are 
made, especially those of gases and vapours, are taught in physics. 
In order to thoroughly understand them, the student should make 
determinations himself in the chemical or physical laboratory. 
The determination of the density of a compound in the gaseous 
state (its vapour density) is of the greatest importance, since this 
density when multiplied by 28-88 gives the molecular weight of the 
compound (p. 51). 

Every substance possesses its own peculiar specific gravity, 
which is always the same, under the same conditions (particularly 
temperature and pressure), however the substance occurs or how¬ 
ever it may have been obtained. A substance which from its pro- 


70 Text-Book of Inorganic Chemistry. 

perties we might take for iron, but which has a different specific 
gravity, say 9 instead of 7-8, is not iron. Numerous compounds 
exist in organic chemistry which not only possess the same per¬ 
centage composition, but are also so similar in their chemical pro¬ 
perties and behaviour that they might be thought to be identical 
if the specific gravity of their gases were not so different from one 
another, which proves that they are polymeric and not identical 
bodies. 

One of the most important of the physical properties of a body 
is its capacity for heat or its specific heat. This must be carefully 
determined by one of the various methods known to physicists. 
The specific heat is particularly important in the case of solids, and 
it has been found that if the solid be an elementary body, the 
product of its specific heat into its atomic weight is a constant 
quantity (about 6). This may be expressed in another way by 
saying that the capacity for heat of all elementary atoms in the 
solid state is the same— i.e. the quantity of heat required to raise 
atomic weights of different solid elements (atomic heat ) is a con¬ 
stant quantity. The law is known after the two French chemists 
who discovered it as Dulong and Petit s law. 1 

Crystalline Shape . 2 

When a substance passes from the liquid or gaseous to the solid 
state, its molecules generally arrange themselves so that regular 
symmetrical forms or crystals are produced. These forms are 
always the same for one and the same substance when produced 
under similar conditions, and even when the conditions producing 
the crystals vary the form is usually the same. 

Common salt, sal-ammoniac, and alum under all conditions 
crystallize in the regular system ; sulphur crystallizes in two forms, 
but only two, accordingly as the crystals are produced from a solu¬ 
tion or from the liquid substance by allowing it to cool. 

The plane surfaces composing the regular shapes of crystals 
are called faces ; the inclination of any two faces to one another is 
called the angle between the faces. It is a remarkable fact that 
although the shape of two crystals of the same substance, prepared 
in the same manner, may apparently be very different from one 

1 Introduced by the Editor. A fuller account of this law is given in the 
Appendix. 

2 This short sketch has been introduced by the Editor. 


Physical and Chemical Properties of Bodies. 71 

another, the angle made by any two similar faces always remains 
unchanged. 

It has been found that the faces of crystals can be arranged 
according to certain imaginary lines supposed to be drawn through 
the crystal and called its axes ; and according to the relative length 
and inclination of these axes all crystals are divided into six sys¬ 
tems : the regular, tetragonal, hexagonal, rhombic, monoclinic, and 
triclinic systems, as they are usually designated. 

The Kegular System. 

In this system there are three axes, which are all equal and all 
at right angles to one another. The fundamental forms of this 



Fig. 17. 


system are : the regular octahedron (fig. 16), the cube (fig. 17), and 
the rhombic dodecahedron (fig. 18). Figs. 19 and 20 illustrate two 



Fig. 19. 


Fig. 18. 


common forms of the octahedron and cube combined; in the 
former the cube predominates, in the latter the octahedron. In 
these forms all possible faces are developed, and such crystals, by 



















72 


Text-Book of Inorganic Chemistry. 

far the most common, are called holohedral. Other forms are, 
however, known in which only one half or one quarter of the 
entire number of faces is developed. Such crystals are called 




Fig. 20. 2i. 

hemihedral and tetrahedral respectively. The hemihedral form 
of the octahedron is the tetrahedron (fig. 21). 

Among the large number of bodies crystallizing in this system, 





the following are common examples : common salt (cubes), alum 
(octahedra and cubes), fluor-spar, garnet, the spinelles, galena, 
zinc-blende (often tetrahedra). 

The Tetragonal System. 

In this system the axes are still all at right angles to one ano¬ 
ther, but one is different in length to the other two, which are 
themselves equal. The first axis is th z primary axis, the other two 
are the secondary axes. The fundamental forms are the tetragonal 























Physical and Chemical Properties of Bodies. 73 


pyramid (fig. 22) and the tetragonal prism. Fig. 23 represents a 
combination of one of the prisms with the pyramid, a form in which 
tin-stone often occurs. 

Common substances which crystallize in this system are : tin¬ 
stone, rutile, potassium ferrocyanide. 

The Hexagonal System. 

This system is distinguished from all the others by the fact that 
it possesses four axes. Three of these axes (secondary) are equal 
in length, lie in the same plane, separated from one another by an 
angle of 6o°, and are all at right angles to the fourth or primary 
axis, of which the length is different. Fig. 24 shows a combination 
of the fundamental pyramid and prism—a common form in rock 
crystal. 

The hemihedral forms of this system are very important—many 
substances (eg. calc-spar) assuming these forms only. Fig. 25 
represents the fundamental rhombohedron , and fig. 26 a scaleno- 
hedron. 




A large number of bodies crystallize in the hexagonal system. 
Among them may be named rock-crystal, apatite, calc-spar, sodium 
nitrate. 

The Rhombic System. 

The axes of this system are all of different lengths, but all are 
at right angles to one another. One axis is chosen as primary axis, 
and of the other two secondary axes the longer is called the macro- 




74 


Text-Book of Inorganic Chemistry. 


diagonal , and the shorter the brachy-diagonal. The two figures 
27 and 28 represent the fundamental pyramid and a combination, 
common in crystals of desmine. 



Fig. 27. 



Fig. 28. 



This is also a large system; some of the more*important sub¬ 
stances crystallizing in it are sulphur, topaz, zinc sulphate. 


The IVIonoclinic System. 


All three axes of this system are of unequal length ; of the 
secondary axes one ( ortho-diagonal ) is perpendicular to the pri¬ 
mary axis, and the other is inclined to it at some angle other than 



90° {klino-diagonal). The fundamental pyramid is shown in fig. 
29, and a combination (gypsum) in fig. 30. 

Of substances belonging to this system may be mentioned: 


























Physical and Chemical Properties of Bodies. 75 

sulphur (when melted and allowed to cool), gypsum, orthoclase, 
cane-sugar. 

The Triclinic System. 

In this system, the most unsymmetrical, all the axes are un¬ 
equal, and they are all inclined to one another at some other angle 
than 90°. The two figures (31 and 32) represent the fundamental 
pyramid and a combination occurring in albite. 

Albite, copper sulphate, axinite, and other substances occur in 
crystals belonging to this system. 

Dimorphism. —Substances which, although chemically iden¬ 
tical, possess the power of crystallizing in two entirely different 
forms are said to be dimorphous — e.g. sulphur (see ante). Some 
substances (eg. titanic oxide) can crystallize in three different 
forms and are called trimorphous. 

Isomorphism .—Two bodies are called isomorphous when they 
crystallize in the same form and possess a similar chemical com¬ 
position. Thus: calc-spar (CaCO s ), dolomite (CaMgC 0 3 ), 
magnesite (MgC 0 3 ), spathic iron ore (FeCo 3 ), calamine (ZnC 0 3 ) 
all crystallize in rhombohedra, having nearly the same angle, and 
a similar chemical composition (M^COg). Many other examples 
might be given of this important and useful law. 1 

Numerous methods may be employed to produce crystals 
from gases, liquids, or amorphous solids. Many gases and liquids 
on becoming solid by cooling assume a crystalline state. The 
violet coloured gas of iodine deposits on cooling dark-grey crystals 
of solid iodine. Water crystallizes— i.e. becomes ice—at o° ; liquid 
carbonic acid at 78°. 

Solid bodies, which can be volatilized unchanged, are mostly 
obtained in a crystalline state by sublimation , or by heating the 
solid and gradually cooling the gas which is evolved. In this 
manner, sulphur, arsenious acid, mercuric chloride, &c., may easily 
be crystallized. Most bodies melt before they sublime, others pass 
directly from the solid into the -gaseous state— i.e. they boil at a 
lower temperature than they melt. 

Solid bodies which can be melted may be easily made to 
crystallize by allowing a portion of the liquid to solidify and then 
pouring off the still liquid portion. Crystals of sulphur, bismuth, 
&c., may be easily obtained in this way. 

The commonest method of preparing crystals is, however, to 


1 See Appendix. 


y6 Text-Book of Inorgatiic Chemistry . 

dissolve the substance in some volatile liquid and then allow it 
to evaporate slowly ; common salt crystallizes in cubes from a 
saturated aqueous solution on evaporation; or a hot saturated 
solution may be allowed to cool slowly, as a liquid usually dis¬ 
solves more of a substance when hot than when cold. In this way 
crystals of nitre, Glauber’s salt, and many other salts, are ob¬ 
tained from hot water, sulphur from hot carbon disulphide, mer¬ 
curic chloride from alcohol, graphite from cast-iron, &c. 

Many metals—copper, silver, &c.—may be deposited from their 
aqueous solutions in a crystalline form by electrolysis. But the 
conditions under which most of the numerous natural crystals of 
the mineral kingdom, particularly those which do not dissolve in 
the ordinary solvents (corund, heavy-spar, and many others), have 
been produced are as yet unknown. 

Solubility .—Nearly all solid bodies dissolve in some liquid ; 
but the manner in which solution takes place may be different. 
Many liquids dissolve solids without producing any chemical 
change—-thus, common salt, nitre, &c., dissolve in water, and the 
dissolved substances then separate out again unchanged on the 
removal of the solvent. This kind of solution may be called 
inechanical solution. 

Other liquids dissolve solids by reason of a chemical action 
which takes place between them. Water dissolves sulphuric an¬ 
hydride, phosphoric anhydride, barium oxide, &c., to form chemical 
compounds with these bodies. In other cases chemical change 
and decomposition may take place, as when nitric acid dissolves 
lead oxide or metallic copper. This kind of solution, depending 
upon chemical change, may be called chemical solution. 

Of mechanical solvents, water is by far the most important; 
then follow alcohol, ether, carbon disulphide, chloroform, benzol, 
&c. The solubility generally increases with the temperature, so that 
liquids at the boiling temperature usually dissolve far more of a 
solid than at the ordinary temperature of the air. A few substances 
are as soluble in the cold as in the warm liquid— e.g. common salt 
in water; and a few are more soluble at the ordinary temperature 
than when the liquid is heated, for example, calcium hydrate in 
water—hence lime-water (a solution of calcium hydrate) becomes 
milky when warmed from the separated calcium hydrate, which, 
however, again dissolves on cooling. 

Many bodies which are insoluble in water dissolve in alcohol, 


Physical and Chemical Properties of Bodies. 77 

ether, &c., and vice versa. Barium chloride and Glauber’s salt are 
soluble in water but insoluble in alcohol and ether, while many 
organic substances (< e.g . stearine) are insoluble in water but solu¬ 
ble in alcohol and ether. 

It is of great importance in determining the identity or differ¬ 
ence of two bodies to know the different quantities dissolved by 
their solvents at various temperature's. Such determinations may 
either be referred to the boiling point of the liquid, to o°, or to the 
mean temperature. 

The Chemical Properties of Bodies. 

To these properties belongs the solubility of a substance in 
different solvents, by which a chemical change is produced. It 
may be known that chemical action has taken place if, on 
evaporating the solvent, a chemical compound of the original 
substance remains behind. No metals are mechanically dissolved 
by any liquid except mercury. Some are acted on by water and 
alcohol and dissolved chemically. Most of the metals dissolve 
chemically in some acid, as nitric, hydrochloric, or sulphuric acid, 
or in aqua regia. Mercury, copper, lead, &c., are dissolved by 
nitric acid and converted into nitrates ; while tin is transformed 
into the insoluble dioxide by nitric acid and dissolved by hy¬ 
drochloric acid. v Gold and platinum do not dissolve in nitric or 
hydrochloric acid alone, but a mixture of the two acids (aqua regia) 
at once converts them into the soluble chlorides. 

In the same manner as the metals, most of the metallic oxides 
are insoluble in water, but form with most of the acids soluble 
salts. It is characteristic of many substances that they are in¬ 
soluble or difficultly soluble in water and other solvents : barium 
sulphate, for example, is neither soluble in water nor in dilute acids. 

In distinguishing between different bodies it is especially 
important to notice whether they possess acid or alkaline proper¬ 
ties, or whether they belong to the class of salts or other indifferent 
bodies. Bodies with acid properties which are soluble may often 
be recognized by their sour taste, but more easily and certainly by 
their action on certain vegetable colouring matters. The blue 
solution of litmus or of violets in water is at once turned red by them. 
Paper which has been dipped in a solution of litmus and then 
dried has a blue colour and is called litmus paper ; when a single 
drop of an acid solution is allowed to fall on a strip of this paper 


78 Text-Book of Inorganic Chemistry. 

the spot is at once coloured red. Bodies which, on the other 
hand, possess alkaline properties render a solution of litmus, made 
faintly red with an acid (or red litmus paper), at once blue again, 
while the solution of violets is turned green. In this sense we 
speak of an acid or alkaline reactio?i. 

But besides the soluble bodies with an acid or alkaline reaction 
a large number of insoluble acids and bases • exists which can 
neither be recognized by their ta£te nor by their reaction. That 
certain of these are acid bodies is proved by the fact that they can 
combine with strong bases to form chemical compounds, from 
which they can be again separated by stronger acids. Silica, for 
example, which is insoluble in water and most acids, and which 
therefore neither tastes sour nor possesses an acid reaction, unites 
easily with caustic potash to form a compound called water-glass, 
from which it is again separated by hydrochloric acid. In the 
same manner, black copper oxide, a body quite insoluble in water, 
is readily dissolved by nearly all acids, and is again thrown down 
in the insoluble state by the action of a stronger base, such as 
caustic soda. 

A salt, whether soluble or insoluble in water, is the substance 
produced by the union of a base with an acid. By this process 
the properties of the acid and base are apparently lost. Most 
salts possess neither an acid nor an alkaline taste ; those which 
are soluble in water have a peculiar so called saline taste, and 
generally do not act upon litmus either red or blue. By the 
union of an acid and a base to form a salt both are said to be 
neutralized . 1 

If we take a little strong solution of caustic potash in a beaker, 
and add to it strong nitric acid, with constant stirring, the liquid 
will become strongly heated, and if we at last add the acid care¬ 
fully drop by drop, we shall ultimately reach a point when the 
liquid neither possesses an alkaline nor an acid reaction. The 
liquid is now neutral, and has neither an alkaline nor an acid, but 
a pure saline taste. It now contains a salt—nitre—which crys¬ 
tallizes out in large prisms on cooling the liquid. 

In the same way, if we neutralize caustic soda with sulphuric 
acid, we obtain Glauber’s salt, which also crystallizes out on cooling. 

Are, it might be asked, the sulphuric acid and the soda, the 
nitric acid and the potash destroyed by this process of neutraliza- 

1 For further remarks on the constitution of salts, see sequel. 


Physteal and Chemical Properties of Bodies. 79 

tion and production of a salt, or is it possible to reproduce the 
original substances from the respective salts ? 

This question has been answered by experiment. The repro¬ 
duction of the original acid and base from an aqueous solution of 
the salt may be done in various ways ; the simplest method is to 
pass an electric current by platinum electrodes through a saturated 
aqueous solution of Glauber s salt, coloured violet by a neutral 
solution of litmus. The salt is then decomposed, and, by inter¬ 
action with the water, the alkali (soda) appears at the negative 
pole, while the acid (sulphuric acid) is set free at the positive pole, 1 
these bodies being recognized by the blue colour of the solution 
around the negative pole and the red colour round the positive. 



In order to make this evident a narrow glass vessel (fig. 33) is 
filled with a solution of the salt, coloured violet by litmus, and 
the solution divided into two equal portions by a piece of card¬ 
board placed across it. In each half of the vessel a piece of 
platinum foil is placed, the ends of both being connected with 
the battery wires as shown. It is sufficient to pass the current for 

1 The first decomposition of the salt (Na 2 S0 4 ) is into sodium (Na 2 ) and 
the radical S0 4 ; but the sodium at once decomposes the water, liberating 
hydrogen and forming sodium hydrate (caustic soda) : thus, Na 2 + 2H 2 0 = 
2 NaOH + H 2 ; these two substances appearing at the negative pole. At the 
same time the S0 4 radical breaks up into S0 3 and oxygen, the former uniting 
with water to form sulphuric acid. The substances liberated at the positive pole 
are therefore sulphuric acid and oxygen — Ed. 





















8 o Text-Book of Inorganic Chemistry. 

a few minutes, in order to clearly perceive that the one half 
becomes blue and the other red. 

By this apparatus it may also be shown that the quantity of 
acid set free by the current from the salt exactly corresponds 
with the quantity of the base so liberated ; that, in other words, the 
quantities of the two substances liberated are in the same propor¬ 
tion as they were combined in the original neutral salt. For if, 
after breaking the current, the cardboard partition is removed and 
the liquid well mixed, the acid and base again unite with one 
another, neutrality is restored, both the blue and red colours being 
changed into the original violet. 


WATER. 

We have learnt that water is not, as was previously thought, a 
simple body, incapable of decomposition, but that it consists of 
oxygen and hydrogen. We have further seen how water may be 
built up from its constituents, and in a constant proportion by 
weight or volume, and how it may be again decomposed into them. 

As a chemical compound, water is full of interest. It occurs in 
nature in all three states of aggregation—solid as ice or snow, liquid 
as ordinary water, and gaseous as water-vapour or steam. 

Liquid water is a tasteless, odourless liquid, and is often, but 
wrongly, thought to be colourless ; in thick layers it has a deep- 
blue colour. This colour may be seen by taking a wide glass tube 
about six feet long, of which the sides are thoroughly blackened, 
and one end closed with a plate of glass, and which is filled 
with chemically pure water. If then the tube is held in a vertical 
position, and a white porcelain basin viewed through it, the latter 
will appear beautifully blue. The true colour of water only appears 
when it is perfectly pure ; the slightest trace of an organic sub¬ 
stance renders the colour a more or less pronounced green or 
yellow. It is the blue colour of water which causes the walls of 
the celebrated blue grotto of Naples and everything contained in 
it to be of a blue colour. The opening to the grotto is almost 
closed by water, and it therefore receives nearly all its light 
through the clear, pure water of the sea. 

Ice also possesses a colour of its own. It is bluish-green, and 
may be well seen in the light passing through the walls of grottoes, 
which are sometimes hewn out of glaciers. 



Water. 


81 


Water becomes solid at o°, and crystallizes in the hexagonal 
system, which may be recognized in the irregularly grouped 
crystals of snow, or in water frozen on the windows in winter. On 
the passage of water from the liquid to the solid state, its volume 
does not diminish, but increases. Ice occupies a larger volume 
than the water from which it is produced, and has, therefore, a 
lower specific gravity (0*917) than water at o°, on which it floats. 
In this respect, water forms an exception to the rule that liquids 
in becoming solid contract. This peculiarity is shared by a few 
other bodies, among the metals by bismuth. 

This apparently unimportant property of water is of the greatest 
value in nature. If ice were specifically heavier than water, the 
rivers and oceans would not become covered with a protecting 
crust of ice, as is now the case, but as soon as ice was produced it 
would sink to the bottom, a fresh layer would be formed on the 
surface, this would again sink, until the whole mass of water were 
converted into a solid block of ice, which the heat of summer 
would only very slightly melt. If this were so, it is no exaggera¬ 
tion to say that animal and vegetable life would be impossible on 
the earth. 

Water possesses its greatest density at 4 0 (point of maximum 
density); if heated above this temperature, or if cooled below it, 
water expands. The expansion of water on freezing, although 
only slight, is the cause of many important phenomena. Rocks, 
the crevices and cavities of which become filled with water, are 
broken up and disintegrated by the freezing of this water in the 
winter. In this way the process is commenced by means of which 
the hard rocks are converted into friable, fruitful soil. The force 
of the expanding ice is so powerful that a hollow iron vessel com¬ 
pletely filled with water is burst when the water freezes. 

Water is a volatile substance, and even evaporates at the 
ordinary temperature of the air; ice is also volatile, and gives off 
vapour at temperatures below o°, without melting. Water boils at 
ioo° and then forms a colourless vapour or gas, commonly called 
steam. This vapour becomes visible on escaping into the air 
owing to its partial condensation. The boiling-point of water, like 
that of all other liquids, depends upon the pressure. It is ioo° 
under the ordinary pressure of the air (760 mm.), and rises to 120° 
under a pressure of two atmospheres, and higher still when the 
pressure is still more increased, as in the boilers of high-pressure 
steam-engines. Under a less pressure than the ordinary atmospheric 

G 


82 


Text-Book of Inorganic Chemistry. 

pressure— e.g. on high mountains, or under the receiver of an air- 
pump—water boils ac a lower temperature than ioo°. 

Water as it occurs in nature is never chemically pure ; it con¬ 
tains varying quantities of mineral and organic substances dissolved 
or suspended in it. Natural waters may be divided according to 
their occurrence into : rain-water, spring-water, river-water, 
mineral waters, and sea-water. The last named contains the 
greatest quantity of dissolved mineral substances, principally 
common salt, and in such quantities as to make it quite unfit for 
drinking purposes. 1 Mineral waters, of which many are used for 
medicinal purposes, contain very different substances dissolved in 
them, often in considerable quantities. We distinguish : saline- 
waters, which contain as their chief constituent common salt, or, 
when bitter, magnesium sulphate (Epsom); carbonated-waters , with 
free carbonic acid and other mineral constituents (Apollinaris, 
Seltzer) ; alkaline-waters , with considerable quantities of sodium 
carbonate, &c. (Vichy); chalybeate-waters , containing ferrous car¬ 
bonate, or some other compound of iron (Spa, &c.); sulphuretted- 
waters, with sulphuretted hydrogen gas dissolved in them (Harro¬ 
gate) ; siliceous-waters, containing silica (geysers of Iceland), &c. 

Ordinary spring- and river-waters contain a smaller quantity of 
mineral constituents than the true mineral waters. The chief 
substances contained in these waters are common salt, calcium 
carbonate, calcium sulphate, magnesium carbonate, magnesium sul¬ 
phate, and free carbonic acid. These substances are derived from 
the rocks of the district in which the spring occurs or through which 
the river flows. Tales sunt aquae qualis terra, per quamfiuunt. I f 
a spring rises, for example, through chalky or limestone rocks, as in 
Kent, Berkshire, &c., the carbonic acid which the water obtains from 
the air enables it to dissolve and take up some of the chalk—calcium 
carbonate. Should the district contain gypsum (calcium sulphate), 
as in Cheshire, the water will dissolve a portion of this substance. If, 
again, the neighbouring rocks contain magnesia (e.g. the dolomite 
rocks of Derbyshire), a portion of the magnesium carbonate will 
pass into solution in the water. Lastly, spring- and river-waters 
may be nearly pure if the rocks through which they pass are 
chemically unacted upon by water. For this reason the natural 
waters in many parts of Wales and Scotland contain remarkably 
small quantities of solid constituents. 


1 Sea-water contains about 3-6 per cent, of solids.—E d. 


Water. 


83 

Rain-water, the purest of natural waters, contains small quan¬ 
tities of carbonic acid, compounds of ammonia and nitric acid, 
particles of dust, &c., which it takes up from the air in falling 
through it. 

Ordinary potable waters are roughly divided into two classes : 
hard and soft. A hard water requires more soap to produce a 
good lather, and does not boil vegetables so well as soft water. 
Hard waters may further be divided into temporarily hard and 
permanently hard waters. The former become soft on boiling or 
on the addition of lime-water ; their hardness is due to calcium 
carbonate, which they contain held in solution by carbonic acid. 
On boiling, the gas is driven off, and the calcium carbonate (chalk) 
deposited ; this is how the fur of kettles and the scale of boilers 
are produced. On the addition of lime-water, the lime combines 
with the carbonic acid to form calcium carbonate, which is pre¬ 
cipitated with that originally present, according to the equation :— 

CO - 0 2 Ca + C 0 2 + Ca(OH) 2 = 2CO-0 2 Ca + H 2 0 . 

Calcium Carbonic Calcium Calcium Water 

carbonate acid hydrate carbonate 

Permanently hard waters contain either calcium sulphate 
(gypsum) or some compound of magnesium. As their name 
implies, they are not softened by boiling, because the salts which 
cause their hardness are slightly soluble in water under all con¬ 
ditions. 

Well-water, especially in towns, or near dunghills or stables, is 
often contaminated with nitre and other nitrates, and organic 
matter, produced from the decomposing animal matter which 
penetrates through the soil. Such water often appears quite bright 
and has no unpleasant taste, but if drunk it may be exceedingly 
injurious to health. Many epidemics of typhoid and other diseases 
may be traced to the use of water contaminated with decomposed 
organic substances. The influence of the drinking water of a town 
on the general health has of late years been more and more gene¬ 
rally acknowledged, and hence the importance of supplying all 
towns with ample quantities of wholesome water. 

Chemically pure water is thus unknown in nature. In order to 
prepare it, ordinary water is converted into steam by heating it, 
and this again condensed in suitable vessels to liquid water. This 
process is called distillation, and the water so prepared is distilled 
water. 


84 Text-Book of Inorganic Chemistry. 

The simplest apparatus for this purpose is an ordinary glass 
retort, furnished with a receiver for cooling, also of glass. For the 
preparation of larger quantities of distilled water the steam is 
evolved from the top of a metallic boiler, best of tinned copper, 
and condensed in a tinned tube by cold water. The process of 
the slow distillation of water goes on largely in nature. Clouds 
are produced from the aqueous vapour formed by the slow evapo¬ 
ration of the water of rivers and the sea ; and rain falls when this 
water vapour is still further condensed. Rain-water or melted snow 
is thus distilled water, and might be used in place of the latter 
if it were not rendered slightly impure by falling through the air. 

In order to determine whether a specimen of water is pure, or 
at least free from mineral substances, a portion may be evaporated 
to dryness in a clean silver or platinum dish. Pure distilled water 
leaves no trace of a solid residue. Should a residue remain which 
blackens on heating, organic matter is present in the water, although 
the water may be contaminated with organic substances and the 
residue not blacken on heating. 

Water is the common solvent for a large number of solids— 
acids, bases, and salts. A solution is said to be saturated, when it 
can take up no more of the substance dissolved in it. The quantity 
of a substance required to produce saturation depends chiefly on 
the temperature ; most bodies are more soluble in hot than in cold 
water, and we therefore distinguish between a hot and a cold satu¬ 
rated solution. 

A hot saturated solution of a substance, which is more soluble 
in hot than cold water, usually deposits the excess on cooling in 
the solid state— often as crystals. 

The hot saturated solutions of some bodies, if allowed to stand 
quietly while cooling, do not deposit the excess of the salt. But 
when the cold liquid is touched with a solid body, or when a 
particle of dust falls into it from the air, crystallization begins 
at once, and in some cases (e.g. sodium sulphate) the quantity of 
salt separated is so great that the whole mass becomes solid. Such 
solutions are called supersaturated solutions*. Even water itself 
possesses a similar property. When in quite clean vessels, free 
from grease, it can be cooled several degrees below o° without 
solidifying; but the slightest shaking is then sufficient to immedi¬ 
ately convert the whole mass into ice. 

Water is also a solvent for many liquids— e.g. sulphuric acid, 


Water . 


85 


alcohol, glycerine. But as the solution in these cases is mutual— 
i.e. water dissolves also in these liquids—we say that sulphuric 
acid and water, or alcohol and water, mix with one another. 
Sometimes the mixture of the two liquids, as those we have named, 
takes place in any proportion, but with other liquids only within 
certain limits. Water and ether mix with one another, but only 
within certain limiting proportions. If an excess of ether is added 
it will remain floating on the surface of the water undissolved. 

Nearly all gases are more or less soluble in water. Some, as 
ammonia and hydrochloric acid, dissolve in very large quantities 
(up to 800 times the volume of the water or more). In these cases 
a large quantity of heat is set free, owing to the change in aggre¬ 
gation, the gas losing its latent heat and becoming liquid. Other 
gases only dissolve slightly in water, as carbonic acid, oxygen, 
&c., while others, again, scarcely dissolve at all— e.g. hydrogen, 
nitrogen. In all cases, when no stable chemical compound has 
been produced, the dissolved gases may be again driven out by 
boiling—for example, carbonic acid and ammonia; hydrochloric 
acid, on the other hand, can be only partially expelled by boiling. 

Many substances, particularly salts, which separate out in the 
crystalline form from aqueous solutions, retain water mostly 
chemically combined in a loose manner. This water is called 
water of crystallization. Crystallized washing soda and Glauber’s 
salt contain more than 50 per cent, of this water. According to 
the temperature at which the crystallization takes place, the 
quantity of water of crystallization taken up by any particular 
compound may vary, but the weight of the water is to the weight 
of the substance with which it is combined always in some 
definite proportion of their molecular weights. In crystallized 
soda, one molecule of sodium carbonate is combined with ten 
molecules of water; common alum contains one molecule of 
potassium-aluminium sulphate united with twelve molecules of 
water. 

Many compounds {e.g. soda) contain their water of crystal¬ 
lization so loosely united, that they part with it, or a portion of it, 
on standing in the air. The crystals then lose their lustre and 
shape, they become dull, opaque, and often fall to pieces, being 
converted into an amorphous powder. This phenomenon is called 
efflorescence. The reverse of this, or the attraction which many 
crystalline and other solids have for water, so that if exposed to the 
air they take up water and ultimately form a liquid, is called deli- 


86 


Text-Book of Inorganic Chemistry. 

qnescence. Calcium chloride and magnesium chloride are ex¬ 
amples. All salts do not combine with water when they crystal¬ 
lize out from an aqueous solution; common salt, nitre, potassium 
chlorate and others crystallize without water. We are unacquainted 
with the cause of this different behaviour. 

Those compounds which do not lose their water of crystallization 
at the ordinary temperature do so at ioo° or a slightly higher 
temperature. Many salts, however, contain a portion, usually one 
molecule, of their water of crystallization more firmly combined. 
Blue vitriol (copper sulphate), which crystallizes with five molecules 
of water, readily gives off four of them at ioo° or 120°, but may be 
heated up to i8o° without losing the last, which it only does at a 
temperature between 18o° and 200°. Green vitriol (ferrous sulphate) 
and other similar salts, which crystallize with seven molecules of 
water, retain their last molecule in the same firm manner. 

This last molecule of water is not really water of crystallization. 
It may be called water of constitution , because it can be displaced 
by salts. Zinc sulphate crystallizes with seven molecules of water 
(S 0 2 * 0 2 Zn + 7 H 2 0 ), but the double salt of zinc sulphate and potas¬ 
sium sulphate, which crystallizes out from a mixed solution of the two 
salts, does not contain seven but only six molecules of real water 
of crystallization (S 0 2 - 0 2 Zn + S 0 2 - 0 2 K 2 + 6 H 2 0 ). The seventh 
molecule of water in the zinc sulphate has been displaced by potas¬ 
sium sulphate. 

If we bring sulphuric anhydride or calcium oxide into contact 
with water, a large quantity of heat is set free and the substances 
unite with one another to form compounds from which the water 
can be only difficultly or not at all expelled by heat. If the pro¬ 
duct of the union of sulphuric anhydride and water, which has the 
composition SO s + H 2 0 , be heated up to 340°— i.e. far over the 
boiling point of water—the water is not expelled, but the whole 
compound distils over. But we can again obtain water from it at 
the ordinary temperature by adding a metallic oxide. In the same 
manner, water is set free on the addition of an acid to the com¬ 
pound of calcium oxide with water. Such compounds are called 
hydrates , and the water, water of hydratio?i (p. 56). 

In reality these hydrates do not contain water as such. The 
compound produced by the union of calcium oxide and water 
(slaked-lime) is not to be considered as a compound of the two 
substances, and cannot, therefore, be represented by the formula 


Water. 


87 


Ca 0 ,H 2 0 . It is a calcium salt, which may be compared with 
calcium nitrate, containing two atoms of hydrogen in the place of 
the two atoms of the radical NO. r Thus :— 


Ca 


(O • N0 0 
jO-NO a 


Calcium nitrate 


Ca 


(OH 

{OH 


Calcium hydrate 


In the same manner, the acid hydrates (the true acids), for ex¬ 
ample, sulphuric acid, nitric acid, are not really compounds of 
water, but salts containing hydrogen instead of a metal 

SO a + H a O = SO.jgg 

The hydrates may be called hydrogen salts , and there are two 
classes to be distinguished : (i.) those in which the hydrogen takes 
the place of a metal—the oxy-acids ; (ii.) those in which the hydro¬ 
gen displaces an acid-radical—the hydrates proper as, calcium 
hydrate, potassium hydrate (KOH), &c. 

Many of these compounds are easily decomposed, some even 
(copper hydrate) under ioo°. Others are so unstable that they 
cannot exist at the ordinary temperature of the air, or under 
ordinary conditions generally. If a strong base is added to an 
aqueous solution of a mercuric salt, mercuric oxide and not mercuric 
hydrate is precipitated ; and from an aqueous solution of a carbon¬ 
ate, a strong acid will liberate carbonic anhydride (commonly 
called carbonic acid), and not its hydrate or the true acid. 


The crystals of many salts, which crystallize without water, 
often enclose small quantities of water mechanically. On heating 
these crystals they break into pieces, often with violence, or, as we 


say, they decrepitate. 

Water is deposited on solid bodies from the air, and the quantity 
depends upon the nature of the substance, temperature, relative 
moisture of the air, &c. In order to remove this moisture, or 
o-enerally to abstract loosely combined water from bodies which 
will not bear heating, the substances are placed under a desiccator— 
i.e. in a closed space in which the air is kept dry, and the water- 
vapour which is given off is chemically absorbed. This may be 
best done by concentrated sulphuric acid. The simplest construc¬ 
tion for a desiccator is a bell-jar with ground edge, which is placed 
on a ground-glass plate over a vessel of sulphuric acid and a separate 
vessel containing the substance. The air in the be jar ecomes 


88 


Text-Book of Inoigamc Chemistry. 

quickly dried by the sulphuric acid, and then the moisture, which is 
evolved from the water into the dry space, is also absorbed. 

The process may be much accelerated, if the air in the bell-jar 
is partially or entirely exhausted. Such an arrangement is shown 
in fig. 34. The ground-glass plate is cemented to a plate of iron, 
which is borne by a brass support. A hole is made in the glass 



Fig. 34- 


and through the iron plate down the brass support to the side tube, 
above which is a stop-cock. The side tube is connected by a thick 
piece of india-rubber tubing with a water-air pump, as shown. If 
now the pump is set in action, the air under the bell-jar soon be. 























Water . 89 

comes partially exhausted, and, by closing the stop-cock, may be 
preserved for several days in this condition. 

In order to drive off hygroscopic water, or water of crystalliza¬ 
tion, from those bodies which can be heated without suffering- 
decomposition, an air-bath is used. Such an apparatus of the 
simplest construction is shown in fig. 35. It consists of a copper 
vessel of which the parts are brazed together and of which the 
front side opens to form a door. The substance to be dried is 
placed on a piece of wire gauze, about 6 centimetres from the bottom 



Fig. 35- 


of the vessel. After closing the door, the vessel is heated by a 
gas-lamp, and the flame regulated by the thermometer inserted 
into the bath from above. In order to get rid of the water-vapour 
which is thus produced, an opening is made in the top near the 
thermometer, and a second in one of the side walls just over the 
bottom. Ry this arrangement, a slow and continual stream of air 
flows through the warm bath, carrying with it the water from the 
substance which is to be dried. 













































90 


Text-Book of Inorganic Chemistry. 


HYDROGEN PEROXIDE. 

Co?nposition : H 2 0 2 . 

Besides water there exists another compound of oxygen and 
hydrogen which contains twice as much oxygen as water. This 
compound, called hydrogen peroxide or hydroxyl, was discovered 
by Thenard in 1818. 

It is a colourless, transparent, viscid liquid, with a peculiar 
odour, mixes with water in every proportion, is not yet known in 
the solid state, and is very easily decomposed. The pure com¬ 
pound is decomposed even at 15 0 or 20° into oxygen and water; in 
an aqueous solution, especially when containing a little hydrochloric 
acid, it is much more stable. 

Hydrogen peroxide cannot be prepared directly from oxygen 
and hydrogen, but is formed when the oxygen of peroxides com¬ 
bines with hydrogen or water. Some peroxides (for example, 
those of manganese or lead) cannot be used for its preparation ; 
it is best to employ barium peroxide, obtained by heating barium 
oxide in a stream of oxygen. 

The barium peroxide, when finely powdered and mixed with 
dilute hydrochloric acid, decomposes, not like manganese peroxide 
when so treated to form water and chlorine, but into barium chloride 
and hydrogen peroxide :— 

Ba 0 2 + 2HCI = BaCl, + H 2 0 2 . 

Barium Hydrochloric Barium. Hydrogen 

peroxide acid chloride peroxide 

The preparation of the pure substance is tedious and difficult. 
The barium may be precipitated out of the solution by dilute sul¬ 
phuric acid, and the hydrochloric acid by silver sulphate, and then 
the sulphuric acid exactly precipitated with baryta water, by which 
means a dilute aqueous solution of the substance is obtained. This, 
on concentrating by evaporation in vacuo, decomposes all the 
more readily into oxygen and water, the more concentrated it be¬ 
comes. 

The chemical behaviour of hydrogen peroxide is very remark¬ 
able. The readiness with which it decomposes into nascent oxygen 
and water makes it one of the most powerful oxidizing agents for 
many substances—^, for arsenic, sulphurous acid, and many 


9i 


Hydrogen Peroxide. 

metallic sulphides. Black lead sulphide is easily converted by it 
into white lead sulphate. In old oil paintings, in which the light 
colours produced by white-lead have gradually darkened owing to 
the formation of lead sulphide, the original colours may be again 
produced by carefully washing with a dilute solution of hydrogen 
peroxide. This then converts the lead sulphide into lead sulphate, 
without damaging the picture. 

Hydrogen peroxide acts, however, upon many substances as a 
reducing agent and abstracts oxygen. If an aqueous solution is 
added to silver oxide (or manganese peroxide), the liquid froths up 
and evolves considerable quantities of oxygen. The atom of oxygen 
which the peroxide contains more than water unites with an atom 
of oxygen from the silver oxide to form a molecule, and the silver 
oxide is at the same time reduced to metallic silver :— 

H 2 0 2 + Ag .,0 = H 3 0 + 0 2 + Ag 2 . 

Hydrogen Silver Water Oxygen Silver 

peroxide oxide 

In the same manner manganese peroxide is reduced to man¬ 
ganous oxide, which, when the liquid contains a free acid, unites 
with it, forming a salt. This remarkable reaction may be used to 
determine roughly whether a given solution contains much or little 
hydrogen peroxide. The solution is poured on a little powdered 
black oxide of manganese contained in a test-tube when a rapid 
evolution of oxygen at once begins, sufficient to show its presence 
by kindling a glowing chip of wood. 

To detect traces of hydrogen peroxide in an aqueous acid liquid, 
a little ether is added, and one drop of a solution of potassium 
bichromate; after shaking, the ether, as it rises to the surface, is 
found to be tinged with a beautiful blue colour, due to some com¬ 
pound of chromium probably containing more oxygen than chromic 
acid. 


OZONE. 

Composition : 0 3 . 

It was long known that on turning the glass plate of an elec¬ 
trical machine a peculiar odour may be noticed, very similar to 
that produced when phosphorus is allowed to oxidize slowly in 
moist air. Careful experiments have shown that the cause of this 
odour is a chemical compound produced by these and other pro- 



92 


Text-Book of Inorganic Chemistry. 

cesses from the oxygen of the air. To this body the name ozone 
has been given. 

Ozone is an allotropic modification of oxygen, and its relation 
to ordinary oxygen is somewhat the same as that of ordinary 
phosphorus to the red amorphous modification of this element. 
In the same manner as ordinary phosphorus has a much more 
powerful affinity for oxygen, sulphur, and other elements than 
the red phosphorus, so the affinity of ozone for most substances 
is very much . greater than that of ordinary oxygen. And, 
again, as common phosphorus is converted into the red variety 
on heating, so also is ozone converted by heat into ordinary 
oxygen. 

Ozone has never yet been prepared chemically pure, and particu¬ 
larly never free from ordinary oxygen. In all attempts to convert 
oxygen into ozone, a considerable quantity, always the greater 
part of the former gas, remains unchanged. Our ideas, therefore, 
of the properties of ozone refer only to a mixture of it with oxygen. 

It is a colourless gas, with a very intense odour, soluble in 
water, to which it imparts its smell; it acts in an irritating manner 
when inspired, and attacks the bronchial tubes in the same manner 
as chlorine. Ozone is condensed to the liquid state more easily 
than oxygen, and then forms a blue liquid of varying tints according 
to the pressure. 

Ozone occurs in atmospheric air, although only in very minute 
quantities. It is produced from the oxygen of the air in places 
where water evaporates quickly, particularly in the neighbourhood 
of the sea and of the salt-works in Germany, where the brine 
evaporates spontaneously. Various organic compounds— eg. oil 
of turpentine, oil of bitter almonds—possess the property, when 
exposed to the light, of converting the oxygen of the air into 
ozone, absorbing it mechanically, and then yielding it up to other 
bodies. 

The partial conversion of oxygen into ozone may be brought 
about by allowing electric sparks to pass through the gas, or by 
allowing phosphorus, half covered with water, to oxidize slowly to 
phosphorous acid in a large vessel. In the latter case the ozone 
produced soon combines chemically with the phosphorous acid. 

The best method to obtain ozone in comparatively speaking 
large quantities is to subject a slow stream of oxygen or atmo¬ 
spheric air to the action of high electric potential without allowing 
sparks to pass through the gas. By this silent discharge of the 


Ozone . 


93 


electricity through the oxygen which then takes place several per 
cent, of the gas is converted into ozone—a very remarkable fact, 



Fig. 36. 


















94 Text-Book of Inorganic Chemistry. 

and one which suggests the question whether other gases might not 
also be changed in the same way. 

The figure on the preceding page (fig. 36) shows an apparatus 
which yields large quantities of ozone, and which may be easily 
made by a skilful glass-blower. Through the centre of the wooden 
lid of the cylinder A A passes a wide glass tube, open at the top 
and closed at the bottom. At the lower end of this tube, near the 
side, a narrow piece of glass tubing is melted, which is bent up¬ 
wards and again at right angles ( a) above the lid. By this tube the 
oxygen enters, and by the tube b it leaves the apparatus when 
partially converted into ozone. 

In order to produce this conversion of oxygen into ozone, a 
second rather smaller tube is placed inside the wide one. This 
tube is also open at the top and closed at the bottom, and the two 
tubes are melted together air-tight at the line cc, in such a manner 
that the gas, entering at a , passes through the narrow annular space 
between the two tubes, and finds an exit at b. The inner tube and 
the wide cylinder are filled to about the same height with dilute sul¬ 
phuric acid (1 of acid to 10 of water) \ in each of them dips a piece 
of platinum foil to which is melted a piece of stout platinum wire. 
Both wires are now attached to a powerful induction coil, and at 
the same time dry oxygen is allowed to enter at a. 

The two platinum plates and the sulphuric acid with which they 
are in contact become strongly electrified, and discharge themselves 
silently, without the production of sparks, through the space 
between the two tubes, by means of which the oxygen is partially 
converted into ozone. 

Ozone is one of the most powerful oxidizing agents. It so 
exceeds ordinary oxygen in this respect that it often oxidizes bodies 
at the ordinary temperature which only combine with oxygen when 
heated. The metal mercury, which may be kept for years in 
the air without any alteration, becomes at once covered with a 
film of the oxide when exposed to the action of ozone. Silver is 
also oxidized and converted into the peroxide by ozone. Sulphu¬ 
retted hydrogen and sulphurous acid are rapidly converted into 
sulphuric acid, and black lead sulphide, which remains unchanged 
for years in ordinary oxygen, is at once changed into white lead 
sulphate. 

In the same manner as chlorine, ozone easily expels iodine 
from its compounds. Iodine, when free but not when combined 
possesses the property of colouring starch paste blue, even when 


Ozone. 


95 


only minute traces of it are present. If, therefore, dilute starch 
paste is mixed with a solution of potassium iodide and strips of 
white paper are moistened with the mixture, the paper will become 
blue when brought into contact with air containing ozone, and the 
reaction will be all the more rapid the greater the quantity of ozone 
present. 

This very delicate test enables us to find out whether air con¬ 
tains ozone. 1 The reaction is expressed by the equation :— 


2KI + 0 3 + H 2 0 


2KOH + I, + O,. 


Potassium Ozone Wa^er 

iodide 


Potassium Iodine Oxygen 

hydrate 


The powerful oxidizing action of ozone makes it an energetic 
bleacher, and if it should later happen that ozone could be pre¬ 
pared in such quantities and as easily as chlorine, it would be very 
generally employed for this purpose. The success of the old 
process of meadow bleaching depends probably upon the action 
of ozone, this substance being produced in small quantities by the 
rapid evaporation of the water contained in the linen by the sun’s 
rays. Even indigo, which is otherwise a very permanent colour, 
cannot withstand the action of ozone. It is soon decolorized and 
converted into a new chemical compound by the action of ozone 
upon it. 

The difference between the molecular composition of ozone 
and ordinary oxygen has been determined by experiment. It was 
found that when electric sparks passed through oxygen, or, better, 
when the gas was subjected to the silent discharge, its volume 
diminished. The ozone which was produced must therefore be 
denser than oxygen. Exact experiments yielded the result that 
by this process three volumes of oxygen were condensed to two 
volumes of ozone, so that ozone must be one and a half times as 
dense as oxygen. And it was further observed that when the 
mixture of ozone and oxygen was again converted into oxygen by 
heating it regained its original volume. 

From these experiments, and from the fact that the molecule of 
oxygen contains * two atoms, it is concluded that the molecule of 
ozone contains three atoms. If this is correct, the powerful 
oxidizing properties of ozone are explained on the supposition that 
its molecule breaks up into a molecule and a free atom of oxygen— 

1 This test does not prove conclusively that ozone is present, as other bodies, 
such as nitrous acid, hydrogen peroxide, also colour the paper blue. Ed. 


96 Text-Book of Inorganic Chemistry. 

0 3 = 0 2 + 0—the latter then exercising the same force of affinity 
as oxygen in the nascent state— e.g. when set free by the electric 
current. 

Ozone, like hydrogen peroxide, is thus a valuable and easily 
obtained source of nascent oxygen. It may possibly assume very 
Considerable importance in medicine and the arts at no distant 
date. 


THE HALOGENS OR SALT-PRODUCERS. 

The four elements chlorine, bromine, iodine, and fluorine are 
grouped together under this name because they form salts by 
direct union with the metals (aXy, sea-salt, and yewdco, I produce). 
These salts are called haloid-salts to distinguish them from the 
oxy- and sulpho-salts. The haloid-salts consist of two elements, 
the halogen element and a metal—for example, sodium chloride, 
NaCl; while the oxy- and sulpho-salts contain at least three 
elements, a metal and a simple or compound radical united to¬ 
gether by oxygen or sulphur—for example, potassium nitrate, 
NCVOK. 

It was thought at one time that every acid and every salt 
must contain oxygen, and it was long imagined that this element 
was present in the haloid-salts and that chlorine, bromine, &c., 
were compounds of oxygen. As, however, all attempts to detect 
oxygen in the halogens, and as salts of the halogens really con¬ 
taining oxygen were discovered later—for example, potassium 
chlorate (C 10 2 -OK)—the halogens are now considered as ele¬ 
ments, as long as we are unable to decompose them. 

The chemical compounds of the halogens, particularly those of 
chlorine, bromine, and iodine, are so similar that they are often 
difficult to distinguish from one another. It is thus all the more 
remarkable that these three elements should possess such very 
different physical properties in the free state. Chlorine is a 
greenish-yellow gas, bromine a dark red-brown liquid, and iodine 
is a dark grey crystalline solid with metallic lustre. 




Chlorine. 


97 


CHLORINE. 

Chemical Symbol : CL —Atomic Weight : 35*5. 

This element does not occur free in nature, but only combined 3 
its chief compound is common salt (sodium chloride), in which it 
is united with the metal sodium. But although common salt is 
one of the most widely distributed chemical compounds, chlorine 
has only been known for about 100 years ) it was discovered in 
1774 by Scheele, one of the discoverers of oxygen. 

Chlorine is a greenish-yellow gas with a powerful and peculiar 
odour, and owes its name to its colour (^Xoopoy, greenish-yellow). 
It possesses a density of 2*45 compared with air as unity. Its 
molecular weight is, therefore, 2-45x28-88 = 71 (p. 51), and its 
molecule, like that of hydrogen, contains two atoms. At a tem¬ 
perature of -40°, or under a pressure of about ten atmospheres at 
the ordinary temperature, it is condensed to a dark yellow liquid, 
with a specific gravity of 1-3. It is dissolved by water in larger 
quantities than oxygen. One volume of water at the ordinary 
temperature absorbs about 2-5 volumes of the gas, but at o° a 
much greater quantity. At the latter temperature it forms a 
crystalline compound with water, having the composition, Cl + 
ioH 2 0. 

If chlorine is breathed, even in small quantities, it attacks the 
organs of respiration, causes inflammation of the bronchial tubes, 
and so produces a very troublesome cough. In larger quantities 
it may produce inflammation of the lungs, followed by death. The 
chemist must therefore exercise caution in experimenting with this 
substance. 

In the ordinary sense of the word, chlorine is not, like hydrogen, 
a combustible body— i.e. if strongly heated in the air it does not 
burn. But if an atmosphere of hydrogen be substituted for air, 
chlorine burns freely, and, conversely, a jet of hydrogen burns in 
chlorine with a pale, livid flame, producing hydrochloric acid. 
From this we conclude that the affinity of chlorine for hydrogen 
is much stronger than for oxygen. Chlorine cannot support the 
combustion of a piece of glowing charcoal or of a chip of wood, as 
carbon and chlorine never combine directly with one another. 

For the preparation of chlorine its hydrogen compound (hydro- 

H 


98 Text-Book of Inorganic Chemistry. 

chloric acid) is almost universally employed. In order to set the 
chlorine free from this compound, it is only necessary to abstract 
and convert the hydrogen into some non-volatile chemical com¬ 
pound. It will at once be remarked that the strong affinity of 
hydrogen for oxygen might be employed, and the chlorine obtained 
according to the equation :— 

2HCI + O = Cl 2 + H 2 0 . 

Hydrochloric Water 

acid 

And, as a matter of fact, considerable quantities of chlorine may be 
obtained by passing a mixture of hydrochloric acid gas and air over 
red-hot pieces of brick (Deacon’s process). But by this method 
the chlorine remains mixed with large quantities of nitrogen, de¬ 
rived from the atmospheric air, and the process, though suitable for 
technical purposes, is quite unfit for the preparation of the pure gas. 

It is better to employ compounds containing oxygen, and par¬ 
ticularly those only loosely combined with the element, such as the 
peroxides. Most of the metallic oxides are so changed by the 
action of hydrochloric acid that the corresponding chloride and 
water are produced, thus :— 

MnO + 2HCI = MnCl., + H 2 0 . 

Manganous Manganous 

oxide chloride 

Some peroxides behave in the same manner— eg. manganese per¬ 
oxide :— 

Mn 0 2 + 4HCI = MnCl 4 + 2 H 2 0 . 

Manganese Manganese 

peroxide tetrachloride 

But these metallic perchlorides are very unstable bodies, and are 
easily decomposed into the lower chloride and water :— 

MnCl 4 = MnCl 2 + Cl 2 . 

Manganese Manganous 

tetrachloride chloride 

From these two equations it is seen how chlorine may be pre¬ 
pared from manganese peroxide and hydrochloric acid, and how 
only one half of the chlorine is set free. 

The preparation of chlorine, according to this method, is best 
carried out by placing a small quantity of the powdered black oxide 
of manganese in a large glass flask, and then adding commercial 
hydrochloric acid until the flask is about one-third full. It is not 
advisable to introduce more than this into the flask, since the mass 
froths up during the evolution of the chlorine. When the black 


Chlorine. 


99 


oxide has been thoroughly mixed with the hydrochloric acid, it is 
gently heated over a gas flame. The reaction begins at the 
ordinary temperature, and only requires a very gentle warmth to 
continue it. The atmospheric air is then gradually expelled by the 
heavier chlorine, and finally the gas itself issues from the tube, 
which passes through the cork in the neck of the flask. 

The chlorine so prepared is always mixed with vapours of 
hydrochloric acid. In order to separate these vapours, and, at the 
same time, to dry the gas or abstract the water-vapour which it 
contains, it is passed through two wash-bottles, of which the first is 



half filled with water, and the second with concentrated sulphuric 
acid. The water absorbs all the hydrochloric acid, together with a 
little chlorine, and the sulphuric acid removes most of the water- 
vapour, which is carried over from the first wash~bottle. 

Chlorine, unlike oxygen and hydrogen, cannot be collected over 
mercury, as it combines chemically with this metal at the ordinary 
temperature. If it is to be collected over water, the water is first 
made luke-warm, as warm water absorbs much less of the gas than 
cold. If a vessel is required full of dry chlorine, the tube delivering 
the gas is led to the bottom of the vessel (fig. 37)- The chlorine 
then gradually displaces the lighter atmospheric air, and the 

W 



























































100 


Text-Book of Inorganic Chemistry. 


yellow coloured gas may be readily perceived as it rises in the jar. 
In order to lessen the communication with the external air, the 
mouth of the jar is loosely closed with a plug of cotton-wool. The 
chlorine is led into the jar until the interior appears uniformly 
coloured, by which time most of the air has been expelled. All 
such operations with chlorine must be conducted in a good draught 
of air, on account of the poisonous nature of the gas. 

In this method of preparing chlorine, sulphuric acid and 
common salt may be used instead of the hydrochloric acid. An 
intimate mixture of one molecular weight of common salt with at 
least two molecular weights of manganese peroxide is prepared, 
and an excess of dilute sulphuric acid poured over the mixture. In 
this way all the chlorine is obtained from the common salt, and 
the end products of the reaction are acid sodium sulphate, man¬ 
ganous sulphate, water, and chlorine :— 

f ATI 

2 NaCl + MnO a + 3 S 0 2 ]q H = 


Sodium 

chloride 


SO - 2 0 2 Mn 

Manganous 

sulphate 


Manganese Sulphuric 

peroxide acid 

+ 2S0,{gga + Cl, + 2H 2 O. 

Acid sodium 
sulphate 


But the preparation of chlorine from commercial hydrochloric 
acid and black oxide of manganese is always to be preferred ; it is 
more rapid and more advantageous, and the bye-product, man¬ 
ganous chloride, may be easily purified. 

Chloride may also be obtained by the electrolysis of a saturated 
aqueous solution of common salt, or of hydrochloric acid, by em¬ 
ploying electrodes of gas-carbon, and placing a bell-jar filled with 
the same liquid over the positive electrode. 

The chemical affinity of chlorine for most of the metals and 
non-metals is very powerful, and in many cases it even exceeds 
that of oxygen. The elements, oxygen, fluorine, nitrogen, and 
carbon, do not, however, combine directly with it, even at the 
highest temperatures. But compounds of these elements (except 
fluorine) with chlorine may be prepared in an indirect manner. 

The affinity of chlorine for hydrogen is particularly powerful. 
Equal volumes of the two gases, mixed in darkness or candle-light, 
unite to form hydrochloric acid when exposed to direct sunlight, or 
when brought into contact with a burning body or an electric 
spark. The quantity of heat which is set free on the union of the 


Chlorine. 


\o\ 

two gases and the consequent expansion is so great that even 
thick vessels are burst with a loud explosion. A mixture of one 
volume of hydrogen and one volume of chlorine when ignited pro¬ 
duces the same effects, and for the same reason, as a mixture of 
two volumes of hydrogen with one volume of oxygen. A jet of 
hydrogen, however, when ignited' in the air and plunged into a 
jar of chlorine, continues to burn quietly with a pale flame, also 
producing hydrochloric acid. 

The affinity of chlorine for hydrogen is under some circum¬ 
stances stronger than that of oxygen for the same element. An 
aqueous solution of chlorine, exposed to the light, gradually loses 
its colour ; the chlorine expels the oxygen from the water and 
hydrochloric acid remains behind. 

Sulphur unites with chlorine at the ordinary temperature, and 
more quickly when heated, to form the red volatile liquid, sulphur 
chloride. 

Phosphorus, which at the ordinary temperature neither ignites 
in air nor in oxygen, catches fire at once in chlorine, forming one 
of the chlorides of phosphorus. 

Antimony, when finely powdered and shaken into a jar of 
chlorine, burns with light and heat to form antimonic chloride. 
Copper, in a finely divided state, as Dutch metal, behaves in the 
same manner, the compound produced being cupric chloride 
(CuCl 2 ), corresponding to that formed when copper burns in 
oxygen, cupric oxide (CuO). 

Sodium burns in a stream of chlorine, and becomes incrusted 
with sodium chloride, which is identical in every respect with the 
naturally occurring common salt. , 

A lighted taper, the wax or fat of which consists essentially of 
carbon and hydrogen, when introduced into a jar of chlorine, con¬ 
tinues to burn, but only feebly and with a dull red light, large 
quantities of soot being at the same time liberated. The hydrogen 
of the wax unites with the chlorine, while the carbon is set free. 
The same phenomena is shown still more strikingly with volatile 
hydrocarbons rich in hydrogen—for example, oil of turpentine. 

Many organic substances are decomposed by chlorine at the 
ordinary temperature, and in diffused daylight, the decomposition 
being often accompanied by a change in colour. If chlorine is led 
into a blue solution of indigo, or into ordinary black ink, or if 
chlorine water is added to either of these liquids, they both be¬ 
come yellow owing to the formation of new chemical compounds. 


102 


Text-Book of Inorganic Chemistry. 

Fresh flowers, and calico coloured with organic colours when 
moistened with chlorine, lose their colours, or, in other words, they 
are bleached. 

We are acquainted with the new chemical compounds produced 
by the action of chlorine on indigo and on ink ; we know that 
indigo-blue is converted into yellow isatin chloride, which can be 
separated and obtained in a crystalline form. But in most other 
cases, particularly with regard to the unknown colouring matter of 
flowers, we are as little acquainted with the products of decomposi¬ 
tion, when bleached by chlorine, as we are with the chemical 
nature of the colours themselves. 

The action of chlorine in these processes often consists in the 
abstraction of one or more atoms of hydrogen, which unite with 
the chlorine to form hydrochloric acid, while a similar number of 
atoms of chlorine takes the place of the displaced hydrogen. This 
process of substitution , as it is called, often takes place in many 
colourless organic compounds ; the well-known colourless phenol, 
(carbolic acid), for example, is easily converted by chlorine into 
monochlorphenol, dichlorphenol, and even trichlorphenol. 

In other cases, chlorine acts as an oxidizing agent. We have 
already seen that chlorine can decompose water in sunlight, 
forming hydrochloric acid and oxygen ; and if at the same time 
substances capable of easy oxidation are present—and to these 
belong many organic colouring matters—they become oxidized by 
the oxygen set free from the water. In most cases, as in the 
conversion of indigo into isatin chloride, both processes go on 
together. 

The bleaching properties of chlorine make this substance o^ 
great practical value. Linen which was previously bleached by 
exposing it to sunlight on green meadows for weeks together, is 
bleached by chlorine in as few minutes as the old method required 
weeks. In the arts, free chlorine is, however, not employed, but a 
compound called bleaching powder or chloride of lime , which 
easily yields up its chlorine, and which is much easier to manage 
and much less dangerous than the free gas. Books or engravings, 
which, in course of time, have become stained, may be rendered 
completely white by placing them for a short time in fresh chlorine 
water. The printing ink, which consists essentially of carbon 
(lamp-black), is unacted upon by the chlorine. Characters in 
ordinary writing ink, which is an organic compound, are at once 
completely bleached. 


Chlorine . 


103 


In all cases in which chlorine or bleaching powder has been 
used for bleaching purposes, care must be taken that the chlorine, 
which adheres mechanically, is afterwards completely removed. 
Linen, calico, or paper when bleached with chlorine, retains chlorine 
in the pores so persistently, that it cannot be got rid of by repeated 
washing with water. Alone this** chlorine would have no bad 
effect, but under the action of light it produces with water hydro¬ 
chloric acid, which destroys the organic tissues, rendering the 
fabric so brittle, that it ultimately falls to pieces. The chlorine 
which is left in the pores of the fabrics, after they have been 
bleached and thoroughly washed, must be removed by chemical 
means. For this purpose substances are chosen which are easily 
oxidized and which themselves exert no injurious action, as calcium 
sulphite or sodium thiosulphate. The former substance is trans¬ 
formed by chlorine in the presence of water into calcium sulphate 
(gypsum), the latter into sulphur and sodium sulphate. A 
substance used for this purpose is called an antichlor. Linen 
bleached in this way is whiter and lasts longer than that bleached 
by the old method of exposing it to the light and air. 

The property which chlorine possesses of decomposing and 
chemically altering organic and particularly organized bodies gives 
it another practical application. Chlorine acts as a most powerful 
disinfectant; it destroys the organic substances which collect in 
the air of closed spaces, particularly in the chambers of patients 
suffering from infectious diseases. Experience has shown that 
the air of sick rooms is purer and healthier when small quantities 
of chlorine are allowed to evaporate from time to time. This is 
best done by placing a little chloride of lime in one or more 
saucers, and moistening it with a little common vinegar. Small 
quantities of chlorine are then gradually set free. 

Chlorine Hydrate : Cl a + ioH 2 0. 

Water, which at the ordinary temperature only absorbs a little 
more than twice its volume of chlorine, takes up large quantities at 
o°, and forms a chemical compound, having the composition shown 
above. This compound is produced, as small crystalline plates of a 
yellow colour, by leading chlorine into ice-cold water, until finally 
the whole liquid becomes semi-solid. The crystals may be filtered 
off at a temperature below o°, and freed from adhering water by 
pressing between blotting paper. At a few degrees above o 
the hydrate is again decomposed into its constituents. This 


104 Text-Book of Inorganic Chemistry . 

property affords an easy method for the preparation cf liquid 
chlorine. The chlorine hydrate is pressed into the end of a thick 
bent glass tube, which stands in a freezing- 
mixture, and after a few pieces of calcium 
chloride have been placed above it, the open 
end c (fig. 38), is fused together by the blow¬ 
pipe. If now the limb c is placed in a freezing 
mixture and the other limb a gently warmed, 
the chlorine which is set free condenses in c by its own pressure 
to a yellow liquid, while the water is retained by the calcium 
chloride. 



Fig. 38. 


HYDROCHLORIC ACID (Muriatic Acid). 

Composition : HC 1 . 

This substance, the only compound of hydrogen and chlorine, 
occurs in small quantities in the gases emitted from volcanoes. 

It is a colourless gas which fumes strongly in the air, possesses 
a suffocating odour and taste, and a strong acid reaction. Its 
specific gravity compared with air as unity is 1 *26, and its mole¬ 
cular weight is therefore 1*26 x 28-88 = 36-39 (p. 51), corresponding 
to the formula given it above. Under a pressure of 25 atmo¬ 
spheres and at a temperature of — 4 0 it is condensed to a colourless 
liquid. Water dissolves large quantities of the gas forming ordi¬ 
nary (aqueous) hydrochloric acid. 

Hydrochloric acid is produced by the direct union of chlorine 
and hydrogen, with a large evolution of heat, when about equal 
volumes of the gases are mixed and exposed to direct sunlight, or 
when raised to the temperature of combustion by an electric spark 
or a flame. If a very thin glass flask (or better, a thin glass bulb) 
is filled in the dark with a mixture of equal volumes of hydrogen 
and chlorine, corked firmly, and then exposed to bright sunshine, 
the moment when the first ray of sunlight strikes the flask or bulb, 
the union of the two gases at once follows, and the flask is burst 
by the force of the loud explosion. The combination takes place 
more slowly in diffused daylight, and is only produced suddenly 
with an explosion when the two gases are mixed in exactly equal 
proportions in the dark. This mixture of exactly equal volumes 
may be obtained by the electrolysis of strong aqueous hydrochloric 



Hydrochloric Acid. 


10$ 


acid by gas-carbon electrodes, and collecting the mixed gases 
together after all the air has been expelled from the liquid and the 
porous carbon, and after the liquid over the positive pole has 
become perfectly saturated with chlorine. 

When one volume of chlorine unites with one volume of hydro¬ 
gen no condensation occurs, as whefi oxygen and hydrogen com¬ 
bine together, but two volumes of hydrochloric acid are produced. 
One volume of the compound gas consists, therefore, of half a volume 
of chlorine and half a volume of hydrogen. From this it follows 
that the molecule of hydrochloric acid consists of } molecule = i 
atom of chlorine, and £ molecule = i atom of hydrogen ; from this 
also its density may be calculated, and will be found to agree with 
experiment, thus :■— 


weighs 2-450 
„ 0-069 


1 vol. chlorine 

1 „ hydrogen . 

2 „ hydrochloric acid 


1-26 



Again, since 2*519 parts by weight of hydrochloric acid contain 
2-45 parts of chlorine and 0-069 P art °f hydrogen, the percentage 
composition of the gas, according to the usual simple process, 
is :— 

Hydrogen . . . . = 274 parts by weight. 

Chlorine.— 97*26 „ 

Hydrochloric acid . . . = ioo-oo „ 

In other words, this small quantity of hydrogen suffices to com¬ 
pletely change the properties of more than 35 times as much 
chlorine on chemical union with it. 

That hydrochloric acid consists of hydrogen and chlorine in 
equal volumes may be proved by its electrolytic decomposition. 
A piece of apparatus by which this experiment may be performed 
is shown in fig. 39. Two glass tubes of equal calibre, which are 
closed at the top by stop-cocks, are placed in a vertical position, 
and communicate by a side-tube with a third vertical tube, some¬ 
what longer, and enlarged at the upper end into a bulb. In the 
lower ends of the two tubes are two cylinders of gas-carbon, which 
are fitted in water-tight by caoutchouc stoppers, and which are 
placed in connexion with the two poles of a galvanic battery. The 
tubes are filled with strong aqueoiis hydrochloric acid (or a mixture 




io6 Text-Book of Inorganic Chemistry. 

of hydrochloric acid and a strong solution of common salt) through 
the bulb-tube, the stop-cocks being open. On commencing the 
experiment, and for some time afterwards, the stop-cocks are left 
open, so that the liquid in the tube where the chlorine is liberated 
may become saturated ; the chlorine which is evolved being carried 
away by a downward draught as shown in the diagram. If finally 



Fig. 39- 


the two stop-cocks are closed at the same time, hydrogen collects 
in one tube and chlorine in the other in equal volumes, and the 
liquid remains during the experiment at equal heights in the two 
tubes. The liquid which is displaced by the gases rises in the 
third tube and collects in the bulb. We see, therefore, from this 







































Hydrochloric Acid. ioj 

experiment that hydrochloric acid contains chlorine and hydrogen 
in equal volumes. 

For the ordinary preparation of hydrochloric acid gas, common 
salt is almost exclusively used. This salt (sodium chloride), when 
mixed with concentrated sulphuric acid, gives a large quantity of 
the gas at the ordinary temperature "Without heating, and is con¬ 
verted into acid sodium sulphate :— 

NaCl + SO a {°g = HC 1 + S 0 2 j°* a 

Sodium Sulphuric Acid sodium 

chloride acid sulphate 

But if the ordinary crystallized salt, which offers a large surface 
to the acid, is employed, the mixture always froths over. It is 
therefore better to use salt which has previously been melted and 
afterwards broken into moderately large pieces. This evolves the 
gas more slowly, as its surface is much smaller. The action com¬ 
mences at the ordinary temperature, but must afterwards be sup¬ 
ported by a gentle heat. 

If rock-salt is used for the preparation of hydrochloric acid, 
the sulphuric acid must be mixed beforehand with an equal volume 
of water, and, to free the gas from, aqueous vapour, it must be 
passed through wash bottles containing concentrated sulphuric 
acid. 


Aqueous Hydrochloric Acid. 

In the ordinary language of the laboratory, hydrochloric acid 
means the solution of the gas in water. This is produced by pass¬ 
ing a stream of the gas through vessels containing water, which 
must be kept cold as long as absorption is going on. One volume 
of water at o° can absorb about 500 volumes of hydrochloric acid 
gas, at 20 0 a less quantity, but still over 400 volumes. During this 
absorption a considerable quantity of heat is set free;, partly in 
consequence of the chemical attraction between the gas and the 
water, and partly because of the change from the gaseous to the 
liquid state, It is therefore necessary to keep the water cool, 
especially if a cold saturated solution of the gas is required. 

This solution of the gas in water (ordinary hydrochloric acid) 
fumes in the air and possesses the odour of the gas. It reacts and 
tastes acid even when very largely diluted with water. 

A completely saturated solution of hydrochloric acid at o° has 
a specific gravity of 1*21. On heating such a solution considerable 


io8 Text-Book of Inorganic Chemistry. 

quantities of the gas are evolved, and the specific gravity dimin¬ 
ishes. When the specific gravity has become 1*145 liquid boils 
at ioo°. Even at this temperature more hydrochloric acid distils 
over than water ; but on continued boiling the relative quantity of 
the water passing over with the acid gradually increases, until the 
boiling point reaches 1 io°, when it remains constant. The liquid 
which distils over at this temperature has a constant composition 
and a constant specific gravity of i*i. 

This acid, which distils over at a constant temperature and 
which has a constant composition, does not fume in the air. It 
contains about 20 per cent, by weight of the gas. It is not, how¬ 
ever, a true chemical compound, as its composition is altered if 
the pressure at which it is made to boil is changed. The satu¬ 
rated solution at o° contains about 42 per cent, by weight of the 

gas- 

impure aqueous hydrochloric acid is obtained in large quan¬ 
tities as a bye-product in alkali works, in that part of the process 
in which sodium chloride is converted into sodium sulphate by 
the action of sulphuric acid. The hydrochloric acid gas, which is 
given off in immense quantities, is led into large earthenware jars 
half filled with water, which communicate with one another, so 
that the gas unabsorbed in the first is absorbed in the second, 
and so on. 

The commercial hydrochloric acid so obtained fumes in the air 
and has generally a specific gravity of about 1*16, which corre¬ 
sponds to about 33 per cent, of hydrochloric acid by weight. It 
is generally coloured yellow owing to the presence of iron (ferric 
chloride) or organic substances, and contains besides small quanti¬ 
ties of sodium chloride and sulphate, and always arsenious chloride 
if the sulphuric acid employed for its production contained arsenic. 
The presence of this substance may be recognized by the yellow 
precipitate of arsenious sulphide, produced when sulphuretted 
hydrogen is added to the acid. 

The commercial acid, free from arsenic, may be employed for 
the preparation of chlorine, and in all cases when the quite pure 
acid is not required. It may also be employed for the preparation 
of the pure non-fuming acid of specific gravity i*i, by distilling in 
a glass retort until the boiling-point has become constant at no° 
and then collecting the distillate. At least one-fourth of the whole 
must be left behind in the retort to retain all the impurities. 

Hydrochloric acid is one of the strong inorganic acids, and 



Oxides and Oxy-acids of Chlorine. 109 

expels many other acids from their salts. With the metallic 
oxides it forms metallic chlorides and water— e.g. 

CuO + 2HCI = CuCl 2 + HoO. 

Cupric Hydrochloric Cupric Water 

oxide acid chloride 

In the same manner it decomposes many metallic sulphides— e.g. 
ferrous sulphide, antimonious sulphide—liberating sulphuretted 
hydrogen. It also serves as a solvent for metals. Tin, which is 
oxidized but not dissolved in nitric acid, easily dissolves in hydro¬ 
chloric acid, forming stannous chloride and hydrogen :— 

Sn + 2HCI = SnCLj + H. 

Hydrochloric Stannqus 

acid chloride 

With most of the peroxides it produces chlorine—^, manganese 
peroxide, lead peroxide. 

The presence of hydrochloric acid or of a soluble chloride in a 
liquid may be recognized by the white precipitate of silver chloride 
produced when silver nitrate is added to the solution. The pre¬ 
cipitate is easily dissolved by ammonia and turns black in the 
light. All chlorides yield chlorine when heated with manganese 
peroxide and sulphuric acid, and the free chlorine may be easily 
recognized by its odour and by its bleaching properties. 


OXIDES AND OXY-ACIDS OF CHLORINE. 

The chemical affinity of chlorine for oxygen is so slight that the 
two elements cannot be made to unite directly with one another. 
But compounds of chlorine and oxygen are produced when the one 
element is brought into contact with the other in the nascent state, 
or by other indirect means. 

We have seen that chlorine and hydrogen unite together only 
in one definite proportion, but chlorine and oxygen form com¬ 
pounds in five different proportions, of which some are only known 
as oxygen compounds, others as oxy-acids, or in combination with 
bases as salts. With the exception of perchloric acid, these com¬ 
pounds are very unstable, and easily decompose into their con¬ 
stituent elements, often with an explosion. 



I 10 


Text-Book of Inorganic Chemistry. 


We distinguish the following compounds of chlorine and oxy¬ 
gen, in which the chlorine atom may have a valency from one to 
seven :— 


Oxides. 


Oxy-acids. 


Hypochlorous anhydride 

. C 1 2 0 

Hypochlorous acid 

j CIO • H 

Chlorous anhydride . 

. CloO, 

Chlorous acid . 

CIO -OH 

Chlorine peroxide 

. C 10 2 

Chloric acid . 

CIO,-OH 



Perchloric acid 

cio 3 -oh 


CHLORIC ACID. 

Composition : C 10 2 • O H. 

We may well commence our description of the compounds of 
chlorine and oxygen with the most important of them—chloric 
acid. The anhydride corresponding to this acid is unknown, and 
the acid itself is only known in an aqueous solution. In order 
to prepare it a solution of potassium chlorate is mixed with 
fluosilicic acid. The insoluble potassium fluosilicate is filtered off 
and the acid liquid neutralized with baryta water; this solution, 
when filtered and evaporated, deposits crystals of barium chlorate. 
An aqueous solution of this salt is made and dilute sulphuric acid 
carefully added, towards the last drop by drop, as long as a pre¬ 
cipitate of barium sulphate is produced. The acid liquid, contain¬ 
ing the chloric acid, is then filtered off and evaporated in a vacuum 
over sulphuric acid; it must not be heated, as it begins to decom¬ 
pose at temperatures over 40°. 

So prepared, it forms a thick acid liquid which has not yet been 
obtained in a crystalline form. 

The salts of the acid—the chlorates—are, however, more stable, 
and of these the potassium salt is the best known and most impor¬ 
tant. They are obtained when chlorine and nascent oxygen are 
brought into contact with an alkali, under certain conditions. 

Potassium chlorate —C 10 2 . 0 K—is produced in considerable 
quantities when chlorine is led into a hot concentrated solution of 
caustic potash. 1 The substances produced are potassium chlorate, 
potassium chloride, and water, of which the potassium chlorate 

1 The delivery tube must be wide or it will become choked by the potassium 
chlorate. 



Chloric Acid. 


111 

crystallizes out first on evaporation owing to its being less soluble 
in water than potassium chloride. The following equation shows 
the reaction :— 

6KOH + 3 C 1 2 = C 1 CVOK + 5KCI + 3 H 8 0 . 

Potassium Potassium Potassium 

hydrate chlorate . chloride 

It is probable that potassium hypochlorite is first produced with 
potassium chloride, and that the former then breaks up into potas¬ 
sium chlorate and potassium chloride :— 

2KOH + Cl 2 = ClOK + KC 1 + H 2 0 . 

Potassium Potassium Potassium 

hydrate hypochlorite chloride 

3 C 10 K = C 10 2 - 0 K + 2 KC 1 . 

Potassium Potassium Potassium 

hypochlorite chlorate chloride 

Potassium chlorate so prepared always contains more or less 
potassium chloride mixed with it, from which it may be easily and 
completely separated by repeated crystallization. The presence of 
potassium chloride mixed with the chlorate may be easily recog¬ 
nized by the white precipitate of silver chloride produced when a 
few drops of silver nitrate are added to a solution of the salt, silver 
chlorate being soluble in water. 

Potassium chlorate crystallizes in thin iridescent plates, without 
water of crystallization. The salt may also be prepared by the 
direct oxidation of potassium chloride. If a concentrated solution 
of potassium chloride is electrolyzed by a powerful current, using 
platinum electrodes, the nascent oxygen which is set free at the 
positive pole oxidizes the potassium chloride to potassium chlorate. 

We have already stated that potassium chlorate is largely 
employed for the preparation of oxygen (p. 12). The chlorine 
and oxygen are so feebly held together that the latter gas is com¬ 
pletely given off on heating, and potassium chloride remains 
behind :— 

2CICVOK = 2KCI + 3 0 2 . 

Potassium Potassium 

chlorate chloride 

In the first stage potassium perchlorate is produced, and in 
such quantities that the method is used for the preparation of this 
salt:— 

2 C 10 2 - 0 K = ClOg-OK + KC 1 + 0 2 . 

Potassium Potassium Potassium 

chlorate perchlorate chloride 


112 Text-Book of Inorganic Chemistry. 

On heating more strongly this salt is also decomposed into 
potassium chloride and oxygen. 

In consequence of the slight affinity of chlorine for oxygen the 
chloric acid of potassium chlorate and other chlorates acts as a 
powerful oxidizer, something in the same manner as the nitric 
acid of the nitrates, A red-hot piece of charcoal, on which a 
little powdered potassium chlorate is sprinkled, burns brilliantly 
by combining with the oxygen of the salt. A mixture of dried 
flowers of sulphur and powdered potassium chlorate not only 
explodes when heated, but also by percussion or even by the heat 
produced when it is rubbed in a mortar. Caution must, therefore, 
be exercised in experimenting with this mixture, especially when 
charcoal powder has been added. A . gunpowder containing 
potassium chlorate instead of nitre—that is, a mixture of charcoal, 
potassium chlorate, and sulphur—is so explosive that fire-arms 
charged with it are burst, and potassium chlorate is, therefore, 
not employed for the manufacture of such explosive mixtures. 

Chloric and hydrochloric acids mutually decompose one an¬ 
other, even in dilute aqueous solutions. A portion of the oxygen 
of the chloric acid oxidizes the hydrogen of the hydrochloric 
acid, and chlorine together with chloric peroxide are given off. 1 
The process goes on when aqueous hydrochloric acid is added to 
a solution of potassium chlorate :— 

2C1CVOK + 4HCI = 2CIO, + Cl 2 + 2KCI + 2H 2 0. 

Potassium Hydrochloric Chloric Potassium 

chlorate acid peroxide chloride 

Concentrated sulphuric acid acts very energetically on potas¬ 
sium chlorate. The chloric acid which is first produced at once 
breaks up into oxygen, and the reddish-brown gas chlorine peroxide, 
which on heating decomposes further into chlorine and oxygen 
with an explosion. 

Chloric acid is a monobasic acid, and all its salts are soluble 
in water. They may be easily recognized by their oxidizing 
properties and by the production of chloric peroxide when heated 
with concentrated sulphuric acid. 

1 This mixture of gases was called euchlorine by Davy.—E d. 



Perchloric Acid. 


1 13 


PERCHLORIC ACID. 

Composition : C 10 3 • O H. 

This compound is a colourless oily liquid of specific gravity 
178, which does not become solid at — 34 0 , fumes in the air, and 
becomes heated when mixed with water. It is prepared by dis¬ 
tilling potassium perchlorate with a large excess of sulphuric acid. 

Potassium perchlorate, as we have just mentioned (p. m), is 
obtained on heating potassium chlorate. The heating is stopped 
after the molten salt has ceased frothing and has become viscid or 
nearly semi-solid. At this point most of the potassium chlorate 
has been converted into potassium chloride and perchlorate, and 
as the latter salt is even more difficultly soluble in water than 
potassium chlorate, it may be easily purified by dissolving in 
hot water and allowing the solution to crystallize. The salt 
requires seventy times its weight of water at the ordinary tempe¬ 
rature to dissolve it, and separates out from its solution in rhombic 
prisms. 

On heating pure potassium perchlorate in a retort with about 
four times its weight of concentrated sulphuric acid, perchloric 
acid distils over as an oily liquid, usually with a yellow colour. 
When mixed with a small quantity of water it becomes heated and 
gives a hydrate of the composition C 10 3 -OH + H 2 0 , crystallizing 
in fine needles. If this hydrate is heated to 1 io° it decomposes 
into perchloric acid and a hydrate with the composition C 10 3 • OH + 
2 H 2 0 , which remains behind in the retort as an oily liquid, boiling 
only at 203°. 

From these facts it will be seen that, although perchloric acid 
contains more oxygen than chloric acid, it is a much more stable 
compound. This is also shown in its behaviour when treated with 
hydrochloric acid. A mixture of potassium perchlorate (or per¬ 
chloric acid) and hydrochloric acid remains unchanged, while one 
of potassium chlorate and the same acid evolves chlorine , and 
chloric peroxide. In this way it may be readily seen whether 
crystalline potassium perchlorate is contaminated with potassium 
chlorate. Even when only very small quantities of the latter salt 
are present, the hydrochloric acid becomes yellow and gives off 
the odour of chlorine. 


1 


114 Text-Book of Inorganic Chemistry. 

Potassium perchlorate when heated breaks up into potassium 
chloride and oxygen, but requires a higher temperature than that 
necessary to decompose potassium chlorate into potassium per¬ 
chlorate and oxygen. It deflagrates when thrown on red-hot 
charcoal in the same manner and for the same reason as potassium 
chlorate. 

The free acid attacks many organic bodies most energetically ; 
it oxidizes them, often with an explosion. 

Perchloric acid is a monobasic acid, and all its salts are soluble 
in water. 


CHLORINE PEROXIDE. 

Composition : C 10 3 . 

This body is only known as an oxide and does not combine with 
water or with bases. 

Chlorine peroxide is a most dangerous substance to experiment 
with, as it easily decomposes with a powerful explosion into its 
constituents. It is produced, together with perchloric acid, when 
small pieces of fused potassium chlorate a.e treated with pure con¬ 
centrated sulphuric acid and the mixture heated not higher than 
40° on a water-bath. It is then evolved as a dark yellow or reddish- 
brown gas, with a powerful odour, and which often explodes even 
at this temperature. In a freezing mixture the gas condenses to a 
red liquid, which boils at about 9 0 . 


CHLOROUS ANHYDRIDE. 

Composition : C 1 > 0 3 . 

CHLOROUS ACID. 

Composition-. CIO-OH. 

Chlorous anhydride is a yellow gas, with an odour similar 
to that of chlorine. It has a specific gravity of 4-07. Like 
chlorine, it attacks the organs of respiration, and bleaches even 




Chlorous Acid. 


ii 5 


more powerfully than the element itself. Water dissolves about 
ten times its own volume of the gas, becoming of a yellow colour. 
This solution contains chlorous acid—CIO-OH, which has, how¬ 
ever, never been prepared pure. It bleaches like the gas and pro¬ 
duces yellow spots on the skin. 

Chlorous anhydride may be e’asily prepared by abstracting 
oxygen from a dilute aqueous solution of chloric acid. Nitric acid 
may be best used to set the chloric acid free from potassium chlo¬ 
rate and arsenious anhydride for the reduction. The following 
equations show the reactions :— 

(i.) cio 2 -ok + N 0 2 -OH = cio 2 -oh + no 2 -ok. 

Potassium Nitric acid Chloric acid Potassium 

chlorate nitrate 

(ii.) 2C10 2 • OH + As 2 0 3 = Cl 2 O a + H a O + As 2 0 5 . 

Chloric acid Arsenious Chlorous Arsenic 

anhydride anhydride anhydride 

For the preparation of the gas 3 parts of arsenious anhydride 
and 12 parts of potassium chlorate, both in the state of a fine 
powder, are mixed, and then 18 parts of pure nitric acid of specific 
gravity 1*33 and 24 parts of water are added. On gently warming 
the liquid it becomes coloured yeliow, and chlorous anhydride is 
evolved in considerable quantities. The gas may be dried by pass¬ 
ing it over calcium chloride, and may then be collected by dis¬ 
placement. At -18° it is condensed to a reddish-brown mobile 
liquid, which boils a few degrees above o°. 

Chlorous anhydride is a very unstable compound. At a little 
above 50° it decomposes with explosion into its constituents. 
Great care must therefore be taken in its preparation, and only 
small quantities should be prepared at once. 

Its aqueous solution combines slowly with bases to form salts— 
the chlorites—which are mostly soluble in water, and which are 
easily decomposed on evaporation into a mixture of chlorate and 
chloride. The lead salt, obtained by mixing a solution of potassium 
chlorite with lead acetate, is deposited as yellow scales. The 
general formula of the chlorites is CIO • OM', and the acid is there¬ 
fore monobasic. 


1 2 



ii 6 Text-Book of Inorganic Chemistry. 


HYPOCHLOROUS ANHYDRIDE. 

Composition : C 1 2 0 . 

This compound is a yellow gas with an odour similar to that of 
chlorine. It has a specific gravity of 2-97, and condenses at -20° to 
a bright red liquid, boiling at about + 20°. When heated it is easily 
decomposed into its constituents, usually with a violent explosion ; 
the liquid decomposes even when shaken. In its preparation great 
care must therefore be exercised. 

The gas is prepared by treating mercuric oxide with chlorine, 
when the reaction goes on as shown in the equation :— 

2 HgO + 2 C 1 2 - HgO,HgCl 2 + C 1 2 0 . 

Mercuric Mercuric Hypochlorous 

oxide oxychloride anhydride 

The ordinary red oxide of mercury, which has been obtained 
by heating the metal in the air, cannot be employed for the reaction, 
but the yellow oxide, which has been prepared by precipitating 
mercuric chloride with caustic soda, and has been carefully washed 
and dried for some time at 300°, must be used. If dry and pure 
chlorine is slowly led over such mercuric oxide contained in a long 
well-cooled tube, it is absorbed and the hypochlorous anhydride 
gas set free. The gas is easily soluble in water, one volume taking 
up about 200 volumes of the gas. This solution may be considered 
to contain 


HYPOCHLOROUS ACID—(ClOH), 

which has, however, never been prepared in the pure state. A 
similar aqueous solution of hypochlorous acid may be also easily 
prepared by shaking up precipitated mercuric oxide with chlorine 
water. The yellow colour of the solution quickly disappears, and 
the mercuric oxychloride which is formed at the same time may 
be separated by filtration. The salts of this monobasic acid—the 
hypochlorites —are also unknown in the pure state, but mixed with 
other bodies they form important substances in the arts and 
manufactures. 

Bleaching powder , which is manufactured in large quantities 
and chiefly used for bleaching purposes, yields a mixed solution of 
calcium chloride and calcium hypochlorite when dissolved in water. 
This bleaching powder, or chloride of lime (not calcium chloride) 


Hypochlorons Acid. 117 

as it is sometimes called, may be considered as a peculiar com¬ 
pound of calcium, oxygen, and chlorine, having the composition 
f ci 

(CaOCl)Cl or CajQQ. It is prepared bypassing dry chlorine 

over dry slaked lime (calcium hydrate). When no more chlorine 
is absorbed, a white powder, which smells faintly like hypochlorous 
acid, remains behind, having the above composition. This com¬ 
pound, when treated with water, is decomposed into calcium hypo¬ 
chlorite and calcium chloride, according to the equation :— 

2(CaOCl)Cl = (C 10 ) 2 Ca + CaCl 2 . 

Bleaching Calcium Calcium 

powder hypochlorite chloride 

But from this mixture pure calcium hypochlorite cannot be sepa¬ 
rated. The weakest acids, even carbonic acid, separate the hypo- 
chlorous acid from this and all other hypochlorites, and the free 
acid then readily breaks up into chlorine and oxygen. On this 
depends the bleaching and oxidizing action of bleaching powder. 
The articles to be bleached are first dipped in a dilute solution of 
the powder in water, and then submitted to the action of a dilute 
acid. The effect produced is the same as if the articles were 
dipped in chlorine water. The reaction takes place according to 
the two equations :— 

(C 10 ) 2 Ca + 2HCI - CaCl 2 + 2CIOH. 

Calcium Calcium Hypochlorous 

hypochlorite chloride acid 

2 C 10 H 4 - 2 HC 1 = 2 H 2 0 + 2Cl a . 

Hypochlorous 

acid 

For a similar reason bleaching powder is well adapted to disinfect 
sick chambers. If a little of the powder placed in a saucer is 
moistened with vinegar (dilute acetic acid), it slowly gives off 
chlorine (p. 103). 

If bleaching powder is digested with water at the ordinary 
temperature and filtered, the alkaline liquid contains, as we have 
seen, calcium hypochlorite together with calcium chloride and un¬ 
changed hydrate. On boiling this liquid the calcium hypochlorite 
is decomposed into chlorate and chloride :— 

3(C10) s Ca = c!o’-oj Ca + 2CaC1 *- 

Calcium Calcium Calcium 

hypochlorite chlorate chloride 

If chlorine gas is led into a dilute solution of potassium or sodium 


118 Text-Book of Inorganic Chemistry. 

hydrate in the cold, a mixture of the corresponding hypochlorite 
and chloride is obtained :— 

2KOH + Cl 2 = KOC 1 + KC 1 + H 2 0 . 

Instead of the caustic alkalies, their carbonates may also be used. 
No evolution of carbonic acid is noticed at first, owing to the for¬ 
mation of an acid carbonate. As soon, however, as the chlorine 
begins to act upon this acid carbonate, carbonic acid is given off. 
These liquids, which are sometimes called Eanx de Javelle , also 
bleach powerfully on the addition of a dilute acid. 

Hypochlorous acid is a powerful oxidizing agent. A solution 
of bleaching powder or of sodium hypochlorite when added to a 
solution of manganous sulphate produces first a white precipitate 
of manganous hypochlorite and hydrate, which rapidly becomes 
oxidized to the brown manganese peroxide. Similarly, with a solu¬ 
tion of lead nitrate lead peroxide is formed. 


BROMINE. 

Chemical Symbol : Br .—Atomic Weight'. 80. 

This element, which was discovered by Balard in 1826, occurs 
in nature combined with sodium and other metals and usually asso¬ 
ciated with sodium chloride. Sea-water and many mineral springs 
contain small quantities of these compounds of bromine. In the 
mineral kingdom it occurs, particularly in South America, as silver 
bromide, and has lately been found combined with potassium in 
the enormous salt deposits of Stassfurt and other places in 
Northern Germany. The bromine which is now brought into 
trade at a low price is chiefly prepared from this last-named 
source. 

Bromine is a dark red-brown liquid, which is nearly opaque 
even in thin layers, and which possesses a specific gravity of 
3-18. It freezes to a crystalline solid, with a metallic lustre, at 
— 7°‘3, and boils at 63°, forming a dark-brown gas. At the ordinary 
temperature it vaporizes rapidly, so that if a few drops are placed 
in a large empty flask, the whole soon becomes filled with the brown 
vapour. Its vapour density is 5*5, according to which it is more than 
twice as heavy as chlorine and 80 times as heavy as hydrogen. 



Bromine. 


119 


Its molecular weight is therefore 5-5 x 28*88 = 159, and its molecule, 
like that of hydrogen, contains two atoms. Bromine, like chlorine, 
dissolves in water, though in rather less quantity; the solution— 
bromine water —has a brownish colour. In the cold below 4 0 
bromine combines chemically with water to form lromi?te hydrate : 
Br 2 + ioH 2 0, which is decomposed at 15 0 , or is rather more stable 
than the corresponding chlorine compound. 

Bromine has an odour resembling that of chlorine, but more 
intense (whence its name from ^pcopos = a stink). It attacks the 
eyes and the mucous membrane -of the respiratory organs even 
more powerfully than chlorine, and great caution must therefore be 
exercised in performing experiments with it. A single drop allowed 
to fall on the skin produces a painful, slowly-healing wound. 

Bromine may be prepared in the same manner as chlorine by 
decomposing hydrobromic acid with manganese peroxide. But as 
hydrobromic acid cannot be so readily obtained as the commercial 
hydrochloric acid, and is much more difficult to prepare, a mixture 
of potassium or sodium bromide with manganese peroxide and 
moderately dilute sulphuric acid answers the purpose better. The 
hydrobromic acid which is at first produced is at once oxidized to 
bromine and water by the manganese peroxide. The distillation 
is best carried out in a tubulated retort fitting air-tight in a receiver, 
the uncondensed bromine vapour being led away by a good draught. 
The reaction is precisely the same as in the preparation of chlorine, 
substituting bromine for chlorine :— 

2KBr + MnO s + oSOjjoH^ 

2S ° 4 ok + sc V°* Mn + 2H *° + Br *- 

The bromine may be purified by shaking with water, and may be 
again freed from this substance by calcium chloride or concentrated 
sulphuric acid. 

The extraction of bromine on the large scale is carried out in 
just the same manner. For this purpose the mother-liquors of 
the mineral- or sea-waters are employed from which most of the 
less soluble chlorides have crystallized out. But as these mother- 
liquors always contain some chlorides, and as the bromides are 
more easily decomposed by sulphuric acid than the chlorides, 
only enough sulphuric acid and manganese peroxide is added to 
the mixture to decompose the bromides. In this way only bromine 
is obtained. The liberation of chlorine is to be avoided, as it com- 


120 


Text-Book of Inorganic Chemistry. 

bines with the bromine to form a chloride, which is difficult to 
separate from the bromine. Small quantities of chlorine may be 
removed from bromine by distilling with potassium bromide, 
thus :— 

BrCl 3 + 3KBr = 3KCI + 2Br a . 

Bromine Potassium Potassium 

chloride bromide chloride 

In its chemical behaviour, bromine is very closely related to 
chlorine. The differences in their chemical action may almost 
all be ascribed to the weaker affinity of bromine in comparison 
with that of chlorine. This difference may be well illustrated in 
their behaviour with regard to hydrogen. Although the affinity of 
bromine for hydrogen is very great, and although a mixture of 
bromine and hydrogen may be made to combine by heating, the 
union of the two substances cannot be produced by the rays of 
the sun, not even the direct rays. 

The weaker affinity of bromine is also shown by the fact that 
chlorine can expel bromine from most of its compounds. The 
non-metals and metals which burn in chlorine combine under 
similar circumstances with bromine also. Its affinity for oxygen 
and carbon is as weak as that of chlorine for these elements. Of 
the bromides, or the compounds so produced, most are soluble in 
water; lead bromide is difficultly soluble, silver bromide quite 
insoluble. 


HYDROBROMIC ACID. 

Composition'. HBr. , 

This compound, like hydrochloric acid, is a colourless gas 
fuming in the air, with a strong acid taste and odour, and very 
soluble in water. Its specific gravity is 271, corresponding to the 
molecular weight: 81. The aqueous solution agrees also in its 
leading properties with aqueous hydrochloric acid ; when com¬ 
pletely saturated it fumes in the air, and gives off a portion of its 
hydrobromic acid gas on heating. 

As previously mentioned, although sunlight cannot cause the 
chemical union of bromine gas and hydrogen, they may be made to 
unite by heating. But it is impossible to prepare pure hydrobromic 
arid by heating a mixture of potassium or sodium bromide with 



I 2 I 


Hydrobromic Acid. 

concentrated sulphuric acid, because a portion of the hydrobromic 
acid which is set free reduces some of the sulphuric acid, forming 
bromine, water, and sulphurous anhydride, thus :— 

2HBr + SO, | oh = Br 2 + 2H 2 0 + S 0 2 . 

Hydrobromic Sulphuric ' Sulphurous 

acid acid anhydride 

The hydrobromic acid prepared in this way is, therefore, always 
coloured red with bromine gas, and is contaminated with sulphurous 
anhydride. 

The pure, colourless gas may be easily prepared by decompos¬ 
ing phosphorous bromide with water, phosphorous acid being at the 
same time produced : — 

PBr s + 3H s O = PHO |q^ + 3HBr. 

Phosphorous Phosphorous Hydrobromic 

bromide ac id ^ Cld 

For this purpose, I part by weight of amorphous phosphorus 
is mixed with 2 parts of water, and then 10 parts of bromine gradu- 



Fig. 40. 

ally added, drop by drop, from a stoppered funnel fitted into the 
cork of the flask containing the phosphorus and water (fig. 40 ). At 
first every drop of bromine gives rise to a powerful action, accom¬ 
panied with a flash of light, but afterwards the union of the phos¬ 
phorus and bromine goes on more quietly. The hydrobromic acid, 
of which the liberation may be completed by gentle heating, is freed 



122 


Text-Book of Inorganic Chemistry. 

from small quantities of bromine by passing it through a (J-tube 
containing fragments of glass and ordinary phosphorus. 

/ Hydrobromic acid is easily and completely decomposed by 
chlorine. If a cylinder of the dry gas is placed mouth to mouth 
with one containing dry chlorine, hydrochloric acid and bromine 
are produced, the latter being recognized by its brown colour. If 
the chlorine is in excess the brown colour disappears again, and 
bromine chloride is produced as a reddish-yellow mobile liquid. 

In the same manner as hydrobromic acid, all bromides are 
also decomposed by chlorine. If chlorine water is added to a 
solution of a bromide, bromine will be set free, and may be easily 
recognized by its yellow or brown colour. If then a drop of carbon 
disulphide is added and the mixture shaken, this substance will 
dissolve the bromine and its colour will become more marked. 
Hydrobromic acid and all soluble bromides give a pale yellow 
precipitate of silver bromide when mixed with silver nitrate. The 
precipitate only dissolves in ammonia with difficulty. 


OXY-ACIDS OF BROMINE. ' 

In the same manner as chlorine, bromine cannot combine 
directly with oxygen. The combination may be brought about in 
an indirect manner, but no compounds of bromine and oxygen 
alone are known. 


Bromic Acid: Br 0 2 - 0 H.—The potassium salt of this acid 
is produced, together with potassium bromide, when bromine is 
dissolved in concentrated caustic potash :— 

6KOH + 6Br = Br 0 2 -OK + 5KBr + 3 H 2 0 . 

Caustic Potassium Potassium 

P otash bromate bromide 

And as the potassium bromate, like potassium chlorate, is much 
less soluble in water than the bromide, it may be easily obtained 
pure by repeated crystallization. When heated, the salt breaks up 
into potassium bromide and oxygen. 

Bromic acid —which may be separated from potassium bromate 
according to the method given on p. iio for chloric acid—is a 
colourless, powerfully acid liquid, which first reddens and then 



Iodine. 123 

bleaches litmus paper. At ioo° it decomposes into bromine and 
oxygen. 


Hypobromous Acid: BrOH, may be obtained by shaking 
bromine water with precipitated mercuric oxide. The aqueous 
solution is a pale yellow liquid, which acts as a powerful oxidizer, 
and which bleaches like hypochlorous acid. 

If bromine is allowed to evaporate at the ordinary temperature 
under a bell-jar with slaked lime, bromide of lime is obtained, a 
substance which resembles chloride of lime, and which also 
possesses bleaching properties. 


IODINE. 

Chemical Symbol : I .—Atomic Weight : 127. 

Up to the commencement of the present century soda was 
exclusively prepared by burning sea-plants, extracting the burnt 
mass (called kelp, or varec) with water, and crystallizing the soda 
from the solution so obtained. The P’rench chemist Courtois, when 
attempting to obtain chlorine from the residues of the above pro¬ 
cess by heating them with manganese peroxide and sulphuric acid, 
noticed a splendid violet vapour, the investigation of which led him 
to the discovery of iodine (1811). 

The chemical nature of this element was, however, first made 
clear by the investigations of Gay-Lussac four years later. It owes 
its name to the colour of its gas (ZobS^s- = violet-coloured). 

Like chlorine and bromine, iodine does not occur free in nature, 
but always in combination with some of the metals. Sodium iodide 
is almost always present in common salt, although in extremely 
small quantities. Even those mineral springs which contain 
sodium chloride, and which are, comparatively speaking, rich in 
sodium iodide, contain such small quantities of iodine that its 
presence can scarcely be detected without concentration. 

Sea-water contains even smaller quantities of sodium iodide, 
and yet this is the source from which by far the greatest proportion 
of the 100 tons of iodine which is consumed annually is obtained. 
But if we were obliged to procure iodine directly from sea-water, 



124 Text-Book of Inorganic Chemistry. 

the process would certainly be an unprofitable one ; we therefore 
extract it in an indirect manner with the aid of organic nature. 
Sea-plants, and particularly sea-weeds, extract the iodine from the 
sea-water and concentrate it in their structures, and it is from these 
plants alone that it is profitable to extract iodine by chemical 
processes. The manufacture is carried on principally on the west 
coasts of Britain (especially Scotland) and France. The sea¬ 
weeds are gathered during low water, dried and burnt, and the 
ashes thus obtained (called kelp ) are rich in sodium iodide. 

Not only the Algae, but also other sea plants, take up iodine 
from the sea, and the same is true of many marine animals—the 
sponges, &c. In this way iodine is conveyed into the bodies, and 
especially into the fat, of many other marine animals, and occurs, 
for example, in ordinary cod-liver oil. In the mineral kingdom 
iodine is found as silver iodide in Peru, Mexico, &c., and as sodium 
iodide in those salt deposits which are due to the evaporation of 
sea-water. So, for example, in rock salt, and particularly in Chili 
saltpetre, or sodium nitrate occurring in Chili. From the mother- 
liquors of the crude sodium nitrate very considerable quantities of 
iodine are now manufactured. 

Iodine is a solid crystalline substance, with a dark grey colour and 
metallic lustre. It neither conducts electricity nor heat; it is 
brittle, and may be easily powdered, and possesses a peculiar, un¬ 
pleasant odour, similar to that of chlorine and bromine, but less 
intense. It melts at 113 0 , forming a dark brown liquid, boils at 
about 200°, and gives off a dark violet-coloured vapour. On 
cooling, the walls of the vessel become covered with innumerable 
small, lustrous crystals of the solid. Its specific gravity is 4-95, and 
that of its vapour 87 compared with air as unity. Its molecular 
weight is thus 87 x 28*88 = 251 ; whence it follows that its mole¬ 
cule, like that of hydrogen, consists of two atoms. 

Although iodine only boils at about 200°, it is so volatile at the 
ordinary temperature that when left exposed to the air it soon loses 
considerably in weight; and when kept in closed vessels it sub¬ 
limes, like camphor, on to the cooler portions of the bottle. It is 
slightly soluble in water, much less than chlorine and bromine, and 
forms a yellow-brown solution. It possesses a strong unpleasant 
taste, colours the skin yellow, and is poisonous. 

Water containing salts dissolved in it, particularly potassium 
iodide, dissolves larger quantities of iodine than pure water ; it is 


Iodine . 


125 


also dissolved in considerable quantities by aqueous hydriodic 
acid. Alcohol and ether, as well as chloroform and carbon di¬ 
sulphide, are good solvents of iodine. Its alcoholic solution pos¬ 
sesses a dark brown colour and is called tincture of iodine. The 
solution in ether is also of a brown colour, while that in chloroform 
or carbon disulphide is dark violet, or when largely diluted a 
bright pink. 

Many applications are made of iodine in medicine and in the 
arts—for example, in photography. Tincture of iodine is applied 
externally to reduce goitre, enlarged joints, and other similar swel¬ 
lings, and potassium iodide is taken internally to produce the same 
effect. Long before this action of the salt was known, even before 
iodine itself was discovered, the ashes obtained by burning sponges 
were used as a medicine for these diseases. It was later found that 
these ashes contain sodium iodide, and since then potassium iodide 
or sodium iodide has always been used in these cases. 

The preparation of iodine from the aqueous extraction of the 
ashes of sea-plants after most of the soda and common salt have 
crystallized out, or from the mother-liquors of Chili saltpetre, 
may be carried out by passing chlorine through the liquid, which 
separates the iodine. Care must be taken that too much of the 
gas is not passed, otherwise iodine chloride will be produced, which 
diminishes the yield. Another method is to heat the crude sub¬ 
stances in iron or earthenware vessels with only just enough sul¬ 
phuric acid and manganese peroxide to decompose the sodium 
iodide present. This salt is more easily decomposed than the 
chloride, and consequently the iodine separates first before the 
chlorine. The process is similar to that used for the preparation 
of chlorine (p. 100), substituting iodine for chlorine 

2NaI + Mn 0 2 + 3S0 2 |q^J = 

2 S 0 2'(ONa + S 0 2 - 0 2 Mn + 2 H 3 0 + I 2 . 

The iodine vapour is received in a number of communicating 
vessels to condense it, then freed from water and purified by re¬ 
sublimation. Usually the two methods are employed together. 
The iodine is first separated by chlorine, excess being avoided, 
again converted into sodium iodide by caustic soda and then dis¬ 
tilled with manganese peroxide and sulphuric acid. Commercial 
iodine not unfrequently contains iodine chloride and cyanide; it 


126 


Text-Book of Inorganic Chemistry. 

may be purified, in the same manner as bromine (p. 120), by distil¬ 
ling with potassium iodide :— 

IC1 3 + 3KI = 3KCI + 2l v 

Iodine Potassium Potassium 

chloride iodide chloride 

In its chemical properties iodine exhibits close analogy with 
chlorine and bromine. It has, however, less affinity for the 
metals and for hydrogen than these two elements, which there¬ 
fore expel iodine from its soluble salts and from its compound 
with hydrogen—hydriodic acid. On the other hand, iodine has a 
stronger affinity for oxygen than the other two halogens ; iodic 
acid, for example, is a solid substance easily crystallized and much 
more stable than the corresponding chloric and bromic acids. 

The behaviour of iodine to starch is very remarkable. Thin 
starch paste mixed with an aqueous solution of iodine becomes of 
a dark blue colour, which, however, vanishes on warming. It 
appears as if a real chemical compound were produced by this 
reaction, but no attempt to determine its composition has been yet 
successful. The colour is only produced by free iodine, and not 
by solutions of its salts, and starch paste can therefore be mixed 
with potassium iodide in any proportion without producing any 
change in colour. But if to the mixture a drop of chlorine water is 
added, which sets iodine free, a dark blue coloration is at once 
noticed. 

This property is employed both to detect traces of starch and 
of free iodine. If a piece of paper moistened with starch is hung 
up in a closed cylinder containing a few drops of a liquid with free 
iodine at the bottom, the paper soon becomes blue owing to the 
vapours of iodine rising from the liquid. 

More delicate even than this reaction, and therefore applicable 
to the detection of infinitesimal quantities of iodine, is the pink or 
purple colour which traces of iodine impart to chloroform or to 
carbon disulphide. Chlorine water is not well adapted for the 
separation of minute quantities of iodine from a liquid, as the slight¬ 
est excess unites at once with the separated iodine to form iodine 
chloride. It is better in such a case to employ fuming nitric acid. 
A few drops of this acid added to the liquid to be tested and well 
shaken up with it in a glass cylinder suffice to convert the iodide 
into a nitrate and set hydriodic acid free ; this at once gives up its 
hydrogen to the loosely combined oxygen of the nitric peroxide 
contained in the fuming nitric acid and sets iodine free :— 


Hydr iodic Acid. 


127 


4HI + N 2 0 4 = 2l 2 + 2H 2 0 + 2NO. 

Hydriodic Nitric Nitric 

acid peroxide oxide 

If a little chloroform or carbon disulphide is then added and 
the liquid well shaken, the former after settling to the bottom of 
the cylinder becomes pink coloured, even if only traces of iodine 
are present. When, on the other hand, iodine is completely absent, 
the chloroform or carbon disulphide remains colourless. 


HYDRIODIC ACID. 
Composition : HI. 


Like hydrochloric and hydrobromic acids, this compound is a 
colourless gas, fuming in the air, with an acid taste and piercing 
odour, and is very easily absorbed by water. In the same manner 
as hydrochloric acid, it consists of one volume hydrogen united 
with one volume iodine vapour to form two volumes of the com¬ 
pound without condensation. Its vapour density (H = 1) is, there- 

fore, ——— = 64, or compared with air as unity 4-4. 

In consequence of the high specific gravity of iodine gas and 
the low specific gravity of hydrogen, the weight of the hydrogen 
contained in hydriodic acid is so small, that it was for some time 
overlooked. From the formulas we see that 127 + 1 = 128 parts by 
weight of hydriodic acid only contain one part by weight of hydro¬ 
gen, or less than one per cent. ; the exact percentage being 


x 100 
128 


= 078. 


The chemical affinity of iodine for hydrogen is even less than 
that of bromine for the same substance, and hydriodic acid cannot 
therefore be obtained by the direct union of its elements. Neither 
can it be prepared pure by decomposing potassium iodide with con¬ 
centrated sulphuric acid. Large quantities of hydriodic acid are, it 
is true, set free by this reaction, but the gas is largely contaminated 
with iodine, because the hydriodic acid is decomposed at the 
moment of its production by the excess of sulphuric acid into 
iodine, water, and sulphurous acid. 

In order to prepare pure hydriodic acid we can only use one 





128 Text-Book of Inorganic Chemistry. 

method—viz. the decomposition of phosphorous iodide by water. 
In order to form this compound ordinary (yellow) phosphorus and 
iodine aie made to combine in the proportion of one equivalent of 
the former to three equivalents of the latter—/.*, in the proportion of 
31 parts of phosphorus to 3 x 127 = 381 parts of iodine, or 1 part to 
12-3 parts. But as it is better to have the phosphorus in slight 
excess, 1 part of phosphorus is mixed with 10 parts of iodine. The 
proper quantity of iodine is weighed off into a dry flask, the 
flask filled with carbonic acid, and then the requisite quantity of 
phosphorus, cut into small pieces and dried, is gradually added. 
The two bodies unite immediately with one another, producing 
flashes of light, to form liquid phosphorous iodide. In order to 
distribute the excess of phosphorus uniformly through the mass, 
the compound is gently heated and kept in a molten state for a 
short time. 

After the compound has cooled and solidified to a crystalline 
mass, it is moistened with a little water and gently warmed. Phos¬ 
phorous acid and hydriodic acid are produced, and the latter 
passes away through a tube fitting in an india-rubber stopper in the 
neck of the flask. It may be led into water if an aqueous solution 
of the gas is to be prepared. The delivery tube must be wide, 
otherwise it is apt to become plugged up for the following reason. 
Phosphorous acid is decomposed on heating into phosphoric acid 
and phosphoretted hydrogen, and this decomposition is greatest, 
the smaller the quantity of water present. But phosphoretted 
hydrogen and hydriodic acid, when in contact with one another, 
unite to form a solid crystalline substance of the composition : 
PHjj'HI = P H, I (phosphonium iodide). In the preparation of 
hydriodic acid this compound may be easily formed, and the 
delivery tube may become plugged with crystals of it. 

Aqueous hydriodic acid is a colourless liquid with a strong acid 
reaction and odour. If completely saturated with the gas in the 
cold, it fumes when exposed to the air. When kept in vessels 
containing air, especially when exposed to the light, the solution 
soon becomes of a yellow or brown colour, from the iodine liberated 
by the oxidation of the hydrogen. The weak affinity with which 
the hydrogen and iodine are combined together is apparent from the 
fact that when a glass rod is heated in the flame, and plunged into 
a jar of the gas, violet vapours of iodine are at once produced. 
If the rod was red-hot, the liberated hydrogen often catches fire. 

Chlorine and bromine decompose hydriodic acid and all the 


129 


Oxy-acids of Iodine. 

iodides (except silver iodide) very easily. If a jar of chlorine is 
inverted over one of hydriodic acid gas, and then the position of 
the jars reversed, so that the heavier hydriodic acid comes .upper¬ 
most, a considerable quantity of heat is produced, and iodine is set 
free, with the production of hydrochloric acid. If the chlorine is 
in excess, the iodine soon disappears, uniting with the chlorine to 
form iodine chloride. 

Hydriodic acid and the soluble iodides produce with silver 
nitrate a yellow precipitate of silver iodide, which is almost inso¬ 
luble in ammonia. 


OXY-ACIDS OF IODINE. 

Of the oxygen compounds of iodine only two are well known- 
iodic acid and periodic acid—which correspond in their composition 
to chloric acid and perchloric acid. 

Iodic Acid : I 0 2 - 0 H. 

This compound is distinguished from the unstable chloric acid 
by its comparative stability. It is a solid, and crystallizes in 
hexagonal plates ; is soluble in water and in alcohol, and possesses 
a bitter acid taste. It may be heated to above ioo° without being 
essentially changed ; at a higher temperature it is decomposed into 
water and iodic anhydride—1 2 0 5 . 

Iodic acid may be prepared by the direct oxidation of iodine by 
nitric acid, if too large quantities are not employed at one time. 
About io grammes of iodine are finely powdered and gently heated 
in a capacious flask with twice as much concentrated nitric acid. 
As soon as no further action takes place the acid liquid is poured 
off, and the residue digested with a fresh quantity of nitric acid 
until all iodine has disappeared. As iodic acid is only slightly 
soluble in nitric acid, the greater portion remains behind in the 
flask as a crystalline powder; the acid liquid is poured off and 
evaporated to dryness, when a further quantity of iodic acid is 
obtained. Both portions are then dissolved in water, evaporated 
to dryness, to expel the last traces of nitric acid, and, finally, heated 
to ioo°—130° in a stream of dry air. 

An aqueous solution of iodic acid does not yield large crystals 
of the acid when evaporated down, but such crystals may be 

K 



130 Text-Book of Inorganic Chemistry. 

easily obtained by adding a little nitric or sulphuric acid, and then 
evaporating. 

Barium iodate is decomposed when boiled with dilute sulphuric 
acid into insoluble barium sulphate and iodic acid. If the clear 
liquid is filtered off, crystals of iodic acid separate on evaporation. 
At the ordinary temperature sulphuric acid has but little action on 
the barium salt. To prepare iodic acid by this process, two parts 
of concentrated sulphuric acid diluted with eight parts of water 
are allowed to act upon nine parts of the finely-powdered barium 
salt. 

If iodine is boiled with caustic potash, the solution contains 
potassium iodide and potassium iodate—just as when chlorine acts 
upon a hot concentrated solution of caustic potash—and the iodate, 
being less soluble than the iodide, crystallizes out first on evapora¬ 
tion. In a solution of potassium iodate, barium chloride produces 
a white precipitate of the difficultly soluble barium iodate. This 
salt may also be directly prepared by boiling iodine with concen¬ 
trated baryta water. 

Iodic acid when heated up to 130°, or slightly higher, loses 
water and becomes converted into a compound having the com¬ 
position I 3 0 7 • OH, which may be considered as one molecule of iodic 
acid united with one molecule of iodic anhydride— i.e. I 0 2 0 H + 
I 0 0 5 . It is distinguished from iodic acid by its insolubility in 
ordinary alcohol. If iodic acid is heated still higher, up to 170°, it 
loses a further quantity of water, and is completely converted into 
the anhydride, I 2 0 5 . The decomposition is also produced by 
absolute alcohol, or by a mixture of this substance with sulphuric 
acid. 


Iodic Anhydride : I 2 0 5 or Jq'| O. 

This compound is easily soluble in water, and then reproduces 
iodic acid. When strongly heated—up to about 300°—it is de¬ 
composed into iodine and oxygen. 

Although the oxygen in iodic acid is much more firmly com¬ 
bined than in chloric or bromic acid, the acid still easily gives up 
its oxygen to those substances which exercise a reducing action, 
and is therefore a powerful oxidizing agent. Sulphurous acid and 
sulphuretted hydrogen, when added to a dilute aqueous solution of 
iodic acid, cause an immediate separation of iodine; nitrous acid, 
phosphorous acid, and other substances produce the same result. 


Oxy-acids of Iodine . 131 

With dry hydrochloric acid gas, dry iodic anhydride yields iodine 
chloride, chlorine, and water :_ 

I 2 0 5 + 10HCI = 2ICI3 + 2 C 1 2 + 5 H 2 0 . 

Iodic acid is a monobasic acid, and produces salts which are 
mostly insoluble in water. It shows a strong tendency to form so- 
called acid salts besides the neutral compounds. The acid salts 
may be considered as neutral salts united with one or two molecules 
of the acid. We know, for example, the following three potassium 
compounds :— 

Normal potassium iodate . . I 0 2 0 K. 

Monacid potassium iodate . . I 0 2 OK + 10 OH. 

Diacid potassium iodate . . I 0 2 OK 4- 2l6 2 OH. 

Like the chlorates, the iodates of the metals are decomposed 
on heating into iodides and oxygen—the iodates of the heavy 
metals usually giving up iodine as well, and becoming converted 
into oxides. It is not, however, possible to expel all the oxygen 
from potassium iodate by heating it even to redness. 

Iodic acid possesses the remarkable property of uniting chemi¬ 
cally with strong acids. If dry and powdered iodic acid is gradu¬ 
ally added to about five times its weight of hot concentrated 
sulphuric acid, a substance separates out on cooling which contains 
both acids chemically united. 


Periodic Acid : IO s -OH + 2 H 2 0 , or IO(OH) 5 . 

The monobasic acid is unknown in the free state ; we are only 
acquainted with its compound, with two molecules of water, which 
may be considered as a pentabasic acid. This substance crys¬ 
tallizes in colourless rhombic prisms, which are soluble in water, 
alcohol, and ether, and deliquesce when exposed to the air. The 
aqueous solution may be boiled without causing decomposition, and 
the crystals are not altered by standing in a desiccator over 
sulphuric acid. At a temperature of 133 0 the acid melts, and 
begins to be decomposed into water, oxygen, and iodic anhydride 
at 140°. 

Four kinds of periodic acid are known in its salts, which are 
mono-, tri-, tetra-, and pentabasic respectively. These acids may 
be represented as compounds of the unknown periodic anhydride : 

k 2 


132 Text-Book of Inorganic Chemistry. 


I 2 0 7 , containing heptad 
thus :— 

i 2 o 7 + h 2 o . . 

I 2 0 7 4 2 H 2 0 . . 

i 2 o 7 + 3 H 2 o . . 
I 2 0 7 + 5 H 2 0 . . 


iodine, with varying proportions of water, 


2 l 0 3 - 0 H. 

oPO^OH), 

°|I 0 2 ( 0 H) 2 . 

2 l 0 2 ( 0 H) 3 . 

2 lO(OH) 5 . 


Of these acids the last-named is known in the free state, but 
the others only in their salts. 

Pentabasic periodic acid may be obtained in the following 
way. When chlorine is led into a hot solution of sodium iodate, 
containing free caustic soda, an acid sodium salt of the pentabasic 
periodic acid is produced according to the equation :— 

I 0 2 - 0 Na + 3 NaOH + Cl 2 = Ioj^Q 1 ^ 3 + 2NaCl. 


This sodium salt, which is nearly insoluble in water, dissolves 
in nitric acid, and gives with silver nitrate a precipitate of the 
corresponding silver salt, which when evaporated down with nitric 
acid, is converted into yellow crystals of the silver compound of 
the monobasic acid : IO s *OAg. Finally, if this salt is boiled with 
water, the former silver compound is again produced, and the 
pentabasic acid set free :— 

2 l 0 3 - 0 Ag + 4 H 2 0 = Io( ( °^ + IO(OH) 5 . 

The solution is then filtered and evaporated down to crystallize. 

The normal silver salt of the pentabasic acid, IO(OAg) 5 , has . 
also been prepared, as well as the sodium salt of the tetrabasic 


acid : O jlo^ONa) 2 ’ by heatin S the above acid salt t0 22 °°~ 

Tn ( (ONa) 3 o {I 0 2 ( 0 Na) 2 HO 

210 1 (OH)' = 0 |l 0 2 ( 0 Na) 2 + 3H *°- 

As will be seen from the above formulae; the composition of the 
periodates is very various and somewhat complex. Similar com¬ 
pounds are known of the oxy-acids of sulphur and phosphorus, 
which have been subjected to more exact investigation than the 


periodates. 







Fluorine. 


33 


Chlorides of Iodine. 

Chlorine and iodine unite together to form two compounds— 
iodine monochloride, and iodine trichloride. 

Iodine monochloride : I Cl, isji red-brown liquid, which crystal¬ 
lizes on standing, and which possesses the odour of both iodine 
and chlorine. The crystals melt at 25 0 ; they are soluble in water, 
whence they may be again extracted by shaking with ether. 

Iodine monochloride is prepared by heating together one part 
of iodine with four parts of potassium chlorate, or by dissolving 
iodine in aqua regia, diluting with water and extracting with ether. 
It then remains behind on allowing the ethereal solution to evapo¬ 
rate. It may also be prepared by acting upon one molecule of 
iodine trichloride with one molecule of iodine :— 

IC 1 3 + I 2 = 3ICI. 

Iodine trichoride\ IC 1 3 , crystallizes as yellow needles when 
chlorine is led over gently heated iodine; if it contains iodine 
monochloride its colour is more or less of a brownish tint. On 
heating, it melts and becomes brown, being decomposed into 
chlorine and iodine monochloride. 


FLUORINE. 

Chemical Symbol : F .—Atomic Weight'. 19* 

This element possesses such powerful chemical affinities for 
other bodies that it has not yet been isolated. It is therefore 
unknown to us in the free state, but the chemical nature of its 
compounds justify us in classing it with chlorine, bromine, and 

iodine. _ . 

Attempts to isolate fluorine have mostly failed, because no 

substance of which vessels might be constructed can withstand its 
action. Glass and porcelain, as well as silver, gold, and platinum, 
are at once attacked by it. 

It is very remarkable that up to the present no compound of 
fluorine with oxygen is known, neither an acid nor a salt. And 
there is, in fact, no element of which we know so few compounds 
as fluorine. It combines neither with chlorine, sulphur, nitrogen, 



134 


Text-Book of Inorganic Chemistry. 

nor carbon, and the only non-metallic elements with which com¬ 
pounds of it are known, are hydrogen, phosphorus, boron, and 
silicon. 

In the mineral kingdom it occurs exclusively in combination 
with various metals, especially with calcium, aluminium, and 
sodium. 

Calcium fluoride is found crystallized in cubes and is called 
fluor-spar. This is by far the most abundant compound of fluorine. 
With aluminium and sodium it forms a double salt having the 
composition Na 3 AlF 3 = 6NaF,Al 2 F 6 . This mineral occurs chiefly 
in Greenland, and is known as cryolite . 


HYDROFLUORIC ACID (Fluoric Acid). 

Composition : HF. 

This compound is a colourless gas with a piercing acid odour, 
and may be condensed by cold to a colourless liquid boiling at 19 0 
and not solidifying at — 34° The gas acts very injuriously upon the 
lungs and air passages when breathed in small quantities, and in 
larger quantities may produce death. Like hydrochloric acid, the 
gas fumes powerfully in the air, and is also energetically absorbed 
by water with considerable evolution of heat. If led into a vessel 
of water which is kept cool, as long as absorption takes place a 
fuming aqueous solution is produced, which gives off hydrofluoric 
acid when heated. At the same time the boiling-point gradually 
rises and at last remains constant (under the normal atmospheric 
pressure) at 120°. The aqueous hydrofluoric acid, which distils 
over at this temperature, contains about 36 per cent, of the gas by 
weight. 

As all siliceous substances are decomposed by hydrofluoric 
acid, with the production of silicon fluoride, the acid cannot be 
prepared in glass or porcelain vessels. It is best to employ 
vessels of platinum, which is not in the least attacked by the acid, 
or of lead, upon which it only slightly acts. A very convenient 
apparatus is made of lead, and has a platinum tube fastened air¬ 
tight into it to conduct the gas into water contained in a platinum 
vessel (fig. 41). 

The vessel which serves as the retort is a cylinder of lead, on 



i35 


Hydrofluoric Acid. 

which fits a cap of the same metal, and which is pierced by a hole 
at the side to receive the platinum tube. Finely powdered fluor¬ 
spar is mixed with concentrated sulphuric acid in the vessel to 
about the consistency of cream, the cap is placed on and made 
tight with plaster of Paris, and, finally, the platinum tube is fixed 
in position and also cemented with plaster. As soon as the plaster 
has set the cylinder is placed on an iron plate and gently warmed. 
The tube is so arranged that it just touches the surface of the 
water which is to absorb the gas, and which is contained in a 
platinum crucible. The crucible must be surrounded with ice or 



cold water in order to keep it cool. By means of this simple 
apparatus, concentrated aqueous hydrofluoric acid may be quickly 
obtained at any time. 

To prepare the gas perfectly pure and absolutely dry it is best 
to employ the compound which is produced by acting on potassium 
fluoride with hydrofluoric acid. This substance, which has the 
composition KHF 2 , decomposes again when heated in a platinum 
vessel into potassium fluoride and hydrofluoric acid. 

Fluorine in combination with hydrogen and with most of the 
metals plays the part of a monad element. But in certain double 
fluorides (e.g. the potassium compound just referred to) it appears 

















136 Text-Book of Inorganic Chemistry. 

as if two atoms of fluorine coalesced to form a dyad, and this 
compound would have a composition analogous to that of potas¬ 
sium hydrate, thus :— 

Potassium hydrate . . . . KO"H. 

Hydric potassium fluoride . . • KF"H. 

By the union of hydrofluoric acid with bases the fluorides are 
produced. Many of the fluorides are soluble in water, and to these 
belongs the silver salt. In this respect hydrofluoric acid differs 
from the otherwise similar acids, hydrochloric, hydrobromic, and 
hydriodic acids, which form insoluble silver salts. On the other 
hand, the fluorides of some metals {eg. calcium), which form very 
soluble chlorides, bromides, and iodides, are quite insoluble in 
water. 

One of the most important properties of hydrofluoric acid, in 
which it is again distinguished from the other halogen acids, is its 
decomposing action on silica and the silicates. No other acid, 
not even sulphuric acid or phosphoric acid, attacks silica, while 
hydrofluoric acid dissolves it, forming gaseous silicon fluoride and 
water:— 

SiO a + 4HF = SiF 4 + 2H a O. 

If the silica is united with bases, as, for example, in common glass 
with lime and soda, the corresponding fluorides are produced by 
the action of the acid. 

We make use of this property not only to test for hydrofluoric 
acid, but also to etch glass—for example, to mark a series of 
divisions on a glass tube. The operation for either of these 
purposes is as follows : Fluor-spar, or the substance to be exam¬ 
ined for fluorine, is finely powdered, placed in an open platinum 
crucible, mixed with concentrated sulphuric acid, and, if necessary, 
very gently warmed. A watch-glass is then taken and its convex 
side covered with a thin layer of wax, by warming it gently and 
rubbing it over with a piece of wax. As soon as the glass is cold 
the characters which are to be etched are written on the waxed 
surface with a sharply-pointed piece of slate-pencil, which removes 
the wax at those points where it touches the glass. If the watch- 
glass so prepared is now placed over the crucible with its convex 
side downwards, the glass becomes corroded at those points which 
are not protected by the layer of wax, and only at those points. 
By the action of the hydrofluoric acid on the glass, silicon fluoride, 
sodium fluoride, and calcium fluoride are produced. The first- 


137 


Hydrofluoric Acid. 

named substance passes away as a gas, the second is removed by 
subsequent washing with water carrying the third mechanically 
with it. 

If the glass is removed after remaining for a short time over 
the crucible, washed with water and the wax removed by warming, 
the characters traced on the glass ‘become distinctly visible, in 
consequence of the contrast between those parts of the glass which 
have been etched and those which were unacted upon by the acid. 

For the production of delicate and exactly equal divisions on 
glass tubes and other similar etchings, it is best to employ the 
commercial aqueous acid, which only fumes slightly in the air. 
The divisions having been marked on the waxed tube, this aqueous 
acid is painted on the parts to be etched with a camel s hair 
brush. 

The fact that hydrofluoric acid etches glass was observed two 
hundred years ago by Schwankhard of Nuremberg— i.e. he found 
that a mixture of fluor-spar and sulphuric acid corroded glass. 
But it was only a hundred years later that Scheele, the discoverer of 
chlorine, showed that this action is due to a gas given off by this 
mixture, and the composition of this gas was first discovered by 
Amp&re at the commencement of the present century. 

Gutta-percha, like wax, is neither attacked nor dissolved by 
hydrofluoric acid, and the aqueous acid can therefore be preserved 
and transmitted from one place to another in bottles of gutta-percha 
provided with stoppers of the same substance. 


ELEMENTS OF THE SULPHUR GROUP. 

To this group belong the three elements, sulphur, selenium, and 
tellurium. They form a natural group similar to those referred to 
on p. 65. All three are solid and crystalline; they can be easily 
melted and sublimed, and are combustible. Selenium and tellu¬ 
rium are distinguished from sulphur by their metallic lustre. 
Their chemical nature is interesting from the fact that they can 
form two distinct classes of compounds in which they play different 
parts. On the one hand, they so closely resemble oxygen that 
they can partially or entirely displace this element from its com¬ 
pounds ; on the other hand, they combine with oxygen itself to 



138 Text-Book of Inorganic Chemistry. 

form compounds usually possessing acid properties, and in which 
they have the same relation to oxygen as phosphorus has in its 
acids and arsenic in its acids. 

In all cases where they take the place of oxygen, their atomicity, 
like that of this element, is always two, while in their compounds 
with oxygen they offer four or six points of attraction to this element. 

When sulphur, which may be here considered as a representa¬ 
tive of the entire group, enters into chemical combination as a 
dyad, the compounds produced—sulphides—exhibit a close che¬ 
mical relationship to the corresponding oxides. Thus :— 


Water. 

. h 2 o 

corresponds to 


Sulphuretted hydrogen 

. h 2 s. 

Arsenious oxide .... 

. As 2 0 3 

corresponds to 


Arsenious sulphide .... 

• As 2 S 3 . 

Potassium hydrate .... 

. KOH 

corresponds to 


Potassium sulphydrate 

. KSH. 

Cupric oxide. 

. CuO 

corresponds to 


Cupric sulphide. 

. CuS. 

Arsenious oxide and arsenious sulpide possess 

acid properties 

while potassium hydrate and sulphydrate, cupric oxide, and sulphide, 

are bases. 


Potassium arsenite is produced by the union of arsenious oxide 

and caustic potash, and, in the same manner, 

from arsenious 

sulphide and potassium sulphydrate potassium sulpharsenite is 

formed :— 


As 2 0 3 + 2KOH = 2As0-0K + 

h 2 o, 

As 2 S 3 + 2KSH = 2 AsS-SK + 

h 2 s. 


So, too, arsenious oxide and cupric oxide unite to form cupric 
arsenite, while the union of arsenious sulphide and cupric sulphide 
produces cupric sulpharsenite; the sulphur compounds contain¬ 
ing in every case sulphur instead of oxygen. 

But when the elements of the sulphur group are united with 
oxygen to form the radicals of the oxy-acids, their chemical 
character is entirely different. In the compounds of dyad sulphur 


139 


Sulphur . 


this element can always be displaced by oxygen ; but this is not 
possible in the second class of compounds, in which the sulphur 
exists as a tetrad or hexad element combined with oxygen, because, 
as far as we know, oxygen only exists as a dyad element, and never 
possesses any higher atomicity. We might imagine the existence of 
a compound corresponding to sulphurous anhydride—S iv 0 2 —in 
which the two atoms of dyad oxygen might be displaced by two 
atoms of dyad sulphur—S iv S u 2 , and the sulphur which is set free 
by the action of sulphuretted hydrogen on sulphurous anhydride is 
perhaps this sulphur :— 

S iv 0 3 + 2H 2 S iS = S iv S u 2 + 2 H 2 0 . 


But we cannot imagine the existence of a similar compound in 
| which the tetrad sulphur of sulphurous anhydride is displaced by 
tetrad oxygen, 0*0“, because, as far as we know, oxygen never 
plays the part of a tetrad. 

The composition of thiosulphuric acid proves that the displace¬ 
ment of oxygen in the oxy-acids, by dyad sulphur, is not only pos- 
j sible, but actually takes place. This acid is sulphuric acid, 


S vi O g 


in which one of the atoms of oxygen united with the 


hydrogen is displaced by an atom of dyad sulphur, thus : 

SvI °4oh- 

These facts are not only true of sulphur compounds, but may 
be generally extended to those of selenium and tellurium. 

Selenium and tellurium may be said to occur as rarely in nature 
as sulphur does abundantly. These two elements (particularly 
selenium) very often occur associated with sulphur. And their 
scarcity is probably the reason why their compounds have not been 
so thoroughlv investigated as those of sulphur. 


SULPHUR. 

Chemical Symbol : S .—Atomic Weight : 32 - 

Sulphur occurs free in nature in considerable quantities ; Sicily 
and other parts of Italy as well as Iceland are particularly rich in 
native sulphur. It is still more widely distributed in chemical 
combination with other bodies. Of the sulphides, the commonest 





140 Text-Book of Inorganic Chemistry. 

are iron pyrites (ferric disulphide, FeS 2 ), copper pyrites (copper-iron 
sulphide, CuFeS 2 ), galena (lead sulphide, PbS), blende (zinc sul¬ 
phide, ZnS), &c. Sulphur is not less common in combination 
with calcium and oxygen, as calcium sulphate, in the mineral 
gypsum (S 0 2 - 0 2 Ca + 2H 3 0), which often occurs in layers of con¬ 
siderable thickness. Other sulphates are also found in nature, 
among which may be mentioned Glauber's salt (sodium sulphate, 
S 0 2 (ONa) 2 , Epsom salt (magnesium sulphate, S 0 2 - 0 2 Mg)—two 
compounds which are contained in mineral waters and in sea 
water. Other naturally occurring sulphates are, heavy-spar 
(barium sulphate, S 0 2 - 0 2 Ba), alunite (basic aluminium-potassium 
sulphate), &c. Besides all these forms in which sulphur is found 
in nature, it also forms a constituent of many important organic 
substances, such as albumen. 

Few substances are so generally known as sulphur, and yet the 
outside world is but little acquainted with its remarkable physical 
and chemical properties. 

Sulphur is a brittle, crystalline solid, with a pale yellow colour, 
without taste or odour, quite insoluble in water, and only slightly 
soluble in alcohol and ether. It is, however, dissolved by carbon 
disulphide in large quantities, especially when the liquid is heated 
to its boiling point. 

At 115 0 it melts to a clear mobile liquid of a pale yellow colour, 
which easily resolidifies to ordinary yellow sulphur, but which if 
more strongly heated (up to 200°) undergoes a remarkable change. 
It does not become more mobile, as might be expected, but more 
and more viscid, and at the same time darker and darker in colour. 
At 200° it is a dark-browm liquid so viscid that the vessel containing 
it may be inverted without any of it running out. When heated 
above 200°, up to 400°, the liquid gradually becomes mobile again, 
without appreciably changing its colour, and boils at 448°. The 
sulphur gas which is so produced and fills the vessel has a dark red- 
brown colour, somewhat resembling gaseous bromine. 

Sulphur which has been heated up to 200° possesses altogether 
other properties than the normal liquid sulphur at 115°. If the 
latter be cooled, for example, by pouring it into cold water, the 
original brittle yellow sulphur, soluble in carbon disulphide, is 
obtained. But if the sulphur at a temperature of 200°, or slightly 
higher, is poured into cold water, a tough, brown, elastic mass is 
produced, which may be drawn out into long threads and which 
only changes into ordinary yellow, brittle sulphur after some time. 


This tough elastic modification of sulphur, unlike ordinary sulphur, 
is insoluble in carbon disulphide. Only after some time, or after 
it has been heated to 111° and allowed to cool again, does it regain 
its solubility in this liquid. 

Sulphur is dimorphous— i.e. it crystallizes in two distinct forms. 
The transparent yellow crystals, as they occur in nature, belong to 
the rhombic system. The two accompanying figures (fig. 42) re¬ 
present two common forms of native sulphur : the one (a) is simple, 
the other (d) has numerous secondary faces. Sulphur crystallized 

in the wet way_ eg. from its solution in carbon disulphide—shows 

exactly the same forms. Large regular crystals of rhombic sulphur 
may be obtained by heating a piece of sulphur with carbon disul¬ 
phide in a sealed glass tube up to 120° or 130° and allowing it to 
cool slowly. 1 The crystals are afterwards freed from the carbon 



b 


a 


Fig. 43 - 


Fig. 421 


disulphide and dried, when they preserve their transparency and 
crystalline form unchanged. 

j On the other hand, when sulphur is melted and quickly cooled, 
it crystallizes in the monoclinic system (fig. 43 )- These crystals 
are best obtained by melting sulphur in a crucible at the lowest 
possible temperature, so that it remains mobile and does not pass 
into the viscid condition, allowing it to cool until a thin solid crust 
forms on the surface of the sulphur, then breaking a hole m this 
crust and pouring out the still liquid sulphur inside If the crucible 
is broken up when cold the interior will be found filled with pale 
yellow needle-like prisms, often as much as an inch long. These 
crystals soon lose their transparency and become opaque and 
brittle * they then consist of a number of minute rhombic crysta s 
similar in form to those in which sulphur occurs in nature, and in 
which it crystallizes from its solution in carbon disulphide. These 

i The glass tube must be thick and well sealed up in the 6lowpipe, other¬ 
wise a very dangerous explosion may be produced. —Ed. 








142 Text-Book of Inorganic Chemistry. 

two modifications of sulphur are also distinguished from one 
another by their different specific gravity. The specific gravity of 
natural or rhombic sulphur is 2*07, while that of the monoclinic 
modification is 1-96. . 

The vapour density of sulphur at different temperatures is 
remarkable. If the molecule of gaseous sulphur consisted, accord¬ 
ing to the general rule, of two atoms, its molecular weight would 

be 2 x 32 = 64, and its density -2*216 (p. 51). But the direct 

28*88 

determination of the density of sulphur vapour at about 500°—a 
temperature considerably above its boiling point—gives quite a 
different number—viz. 6*654, or about three times 2*216. Further 
experiments undertaken to explain this anomaly showed that 
between the temperatures 440° and 850° sulphur vapour did not 
expand regularly with the increase of temperature, and that at 900° 
to i,ooo° its density was only one-third of that at 500°, or about 
2*216. 

From this it appears that there are two polymeric modifica¬ 
tions of sulphur in the gaseous state, one with the density 2*216, 
and the other with a density three times as great—viz. 6*654. The 
molecule in the one case consists, therefore, of two atoms, and in 
the other case of six atoms of the element, and the latter molecule 
breaks up, when heated, into three new molecules, each containing 
two atoms of sulphur. 

Most of the sulphur which comes into trade is brought from 
Italy, and especially from Sicily, where the crude sulphur is partly 
melted out, and partly distilled in cast-iron vessels, from the earthy 
impurities. This sulphur is, however, still impure, and must be 
purified by distillation. For this purpose the arrangement shown 
in fig. 44 is employed. The sulphur is heated to boiling in the 
cast-iron retort G, which is built into the masonry, and of which 
the neck D opens into a large empty bricked chamber A. In the 
same manner as water vapour when cooled at once below its 
freezing-point condenses in the form of snow, so the sulphur vapour 
when issuing into the cold chamber at a temperature below its 
melting point condenses to sulphur snow, which collects on the 
walls and floor of the chamber. This light pale yellow sulphur- 
snow is the flowers of sulphur of commerce. But when the hot 
sulphur vapour has raised the temperature of the chamber to 
the melting-point of sulphur (115 0 ) the sulphur melts and 
collects on the floor as a thin liquid. It may then be allowed 


143 


Sulphur . 


to flow out by an opening which can be regulated from the 
outside. This sulphur is nearly pure. Stick sulphur is easily 
obtained from it by running the liquid sulphur into wooden moulds 
(fig. 45). The sulphur gas which streams into the chamber is at 
a higher temperature than that necessary for the combustion of 
sulphur; it therefore unites with the.oxygen contained in the air 
of the chamber and burns to form sulphurous anhydride, evolving 
at the same time a considerable quantity of heat. In consequence 




Fig. 45 * 


of this increase in temperature the whole of the air of the chamber 
suddenly expands, and the walls would not be able to withstand the 
increased pressure if the chamber were not furnished with a valve. 
This valve opens during the combustion of the sulphur and closes 
afterwards by its own weight. 

The sulphurous anhydride which is now contained in the 
chamber is absorbed by the flowers of sulphur, and if this sulphur 
is afterwards brought into contact with moist air the sulphurous 








































144 Text-Book of Inorganic Chemistry. 

anhydride is gradually oxidized to sulphuric acid. The sulphuric 
acid adheres so persistently to the flowers of sulphur that, even 
after repeatedly washing with water, sulphuric acid may be always 
recognized in the wash-water. For this reason flowers of sulphur 
moistened with water always redden blue litmus paper, which is 
not the case with stick sulphur. 

Sulphur is also prepared by heating iron pyrites in a small 
supply of air, and remelting the sulphur, but the product so ob¬ 
tained is nearly always contaminated with arsenic. 

All elements which unite with oxygen, either more or less 
energetically, also combine with sulphur and in nearly the same 
degree. And sulphur can also unite with oxygen itself. At the 
ordinary temperature the two elements do not unite with one 
another ; but if sulphur is heated in the air, or in oxygen, it catches 
fire and burns with a pale blue flame, to form sulphurous anhydride. 
Besides this body there are several other compounds of sulphur 
and oxygen, of which sulphuric acid is the best known. 

Sulphur unites directly with hydrogen, only with difficulty, and 
the compound of the two elements—sulphuretted hydrogen—is easily 
decomposed again. Chlorine and sulphur combine with one 
another when gently heated. Phosphorus unites with sulphur as 
readily as it does with oxygen, and two of the compounds produced 
correspond in composition to phosphorus and phosphoric anhy¬ 
drides. Nitrogen and sulphur have only a weak affinity for one 
another, and only combine under peculiar circumstances. Finally, 
charcoal (carbon) burns in sulphur gas as it does in oxygen, 
but requires a higher temperature. The compound produced, 
carbon disulphide, has a similar composition to the oxygen com¬ 
pound, carbonic anhydride. 

All those metals which combine with oxygen either at the 
ordinary temperature or when heated burn almost as easily in 
sulphur gas, forming sulphides, with a similar composition to the 
corresponding oxides. Thin copper foil brought into the vapour 
of boiling sulphur burns brilliantly, with a considerable evolution of 
light and heat, to form molten copper sulphide. 

The sulphides of those metals of which the oxides are decom¬ 
posed by heating— e.g. of silver, gold, platinum—are also reduced 
when heated, especially if exposed to the air. Mercuric sulphide 
(cinnabar), which simply volatilizes when heated, forms an excep¬ 
tion to this rule, since the corresponding oxide is decomposed 
into mercury and oxygen by heat. 


vS ulphu retted Hydrogen . 145 

Accordingly as the oxides of the metals are soluble or insoluble 
in water, so are the corresponding sulphides. Soluble potassium 
sulphydrate corresponds to soluble potassium hydrate, calcium 
sulphide to calcium oxide, and insoluble lead sulphide to insoluble 
lead oxide. 

The metallic sulphides which are soluble in water possess the 
property of uniting chemically with more atoms of sulphur than 
are contained in the normal compounds. Potassium sulphide 
(K 2 S), for example, can unite with four more atoms of sulphur, pro¬ 
ducing potassium pentasulphide, which is also soluble in water 

K 2 S + 4S = K 2 S 5 . 

If a solution of this or a similar polysulphide, soluble in water, is 
mixed with hydrochloric acid, the corresponding chloride is formed, 
sulphuretted hydrogen is liberated, and the excess of sulphur sepa¬ 
rates out as a fine, white, amorphous powder. Milk of sulphur is 
this white precipitate which has been washed and dried. The 
equation illustrating its production is as follows :— 

K 2 S. + 2HCI = 2KCI + H 2 S + 4S. 


SULPHURETTED HYDROGEN. 

Composition : H 2 S. 

This compound is analogous to water in its composition, and 
also resembles water in many of its chemical, properties. It is a 
colourless gas with a disgusting odour resembling rotten eggs, has 
a specific gravity of 1-19, and may be condensed, under pressure, 
to a mobile liquid, which solidifies at about - 85° The gas is some¬ 
what soluble in water, to which it imparts its odour; the aqueous 
solution reacting slightly acid with litmus paper. A saturated 
solution in water at the ordinary temperature contains about three 
times its volume of the gas. Alcohol dissolves more, up to five 
times its volume. 

Sulphuretted hydrogen, when inhaled in considerable quantities, 
is poisonous. The chemist, to whom the gas is indispensable, 
becomes more sensitive to the poison the more frequently it is 
inhaled, and the same is also true of other poisons, such as hydro¬ 
cyanic (prussic) acid. In a high state of dilution, sulphuretted 



146 Text-Book of Inorganic Chemistry. 

hydrogen is present in many mineral waters—the so-called sul¬ 
phureous springs (eg. those of Aix-la-Chapelle, Harrogate). These 
waters are often used for medicinal purposes. 

Sulphur and hydrogen unite directly with one another when 
hydrogen is led over heated sulphur ; but as the affinity between 
the two substances is small, only a small quantity of the compound 
is produced. It is better to bring hydrogen in the nascent state 
into contact with sulphur, as by decomposing a suitable metallic 
sulphide with a strong acid. 

The gas is easily obtained in quantity when pieces of ferrous 
sulphide contained in a WoulfPs bottle are acted upon by dilute 
sulphuric or hydrochloric acid. In order to free the gas from 
mechanical impurities it is led through a wash-bottle containing 
water, and is then passed through a calcium chloride tube to dry it. 
It cannot well be collected over mercury, as it is partially decom¬ 
posed by the metal, forming mercuric sulphide. It is best to em¬ 
ploy warm boiled water, which dissolves much less of the gas than 
water at the ordinary temperature. 

The ferrous sulphide used for the preparation of sulphuretted 
hydrogen usually contains some free iron, and the gas is therefore 
generally mixed with a small quantity of free hydrogen. In most 
cases this is immaterial, but if the gas is wanted quite pure it may 
be obtained by acting on antimonous sulphide (black antimony) 
with dry hydrochloric acid gas. The reactions in each of these 
cases are expressed in the following equations : 


FeS + 

Ferrous 

sulphide 

2 HC 1 

= h 2 s + 

FeCl 2 . 

Ferrous 

chloride 

Sb 2 S 3 + 

Antimonous 

sulphide 

6 HC 1 

= 3 H 2 S + 

2 SbCl 3 . 

Antimonous 

chloride 


Sulphuretted hydrogen is combustible, and is easily ignited. If 
a burning body is brought near a jet of the gas it catches fire and 
burns with a pale, bluish flame, producing water and sulphurous 
anhydride. If the supply of oxygen is insufficient, the hydrogen 
only burns, and yellow sulphur is separated. The combustion is 
always accompanied with the separation of sulphur, and it is there¬ 
fore probable that the gas is first decomposed and its constituents 
afterwards burnt. 

In fact, the two elements in sulphuretted hydrogen are only 
very loosely combined. The compound is not only decomposed 


Sulphuretted Hydrogen. 147 

by a high temperature, but also by all oxidizing agents, even the 
weakest, sulphur, being usually separated. Even the oxygen of the 
atmosphere decomposes it in presence of moisture. In a bottle 
half filled with sulphuretted hydrogen water, especially when often 
opened, the odour of the gas gradually disappears and the liquid 
becomes milky from separated sulphur. Sulphuretted hydrogen 
led through water containing iodine in suspension is rapidly decom¬ 
posed, sulphur being set free and hydriodic acid formed ; and this 
reaction may be utilized to prepare a dilute aqueous solution of 
hydriodic acid. Bromine and chlorine act in the same manner, 
but more energetically. Finally, ferric chloride is reduced by 
sulphuretted hydrogen to ferrous chloride with the formation of 
hydrochloric acid and the separation of sulphur :— 

Fe 2 Cl 0 + H 2 S = 2HCI + 2FeCl 0 + S. 

*> rr ic Ferrous 

chloride chloride 

The insolubility of the majority of the metallic sulphides in 
water and the insolubility of some of them in dilute acids makes 
sulphuretted hydrogen a valuable reagent for the precipitation of 
many of the metals from the solutions of their salts. If the gas is 
led through a solution of silver nitrate, or if sulphuretted hydrogen 
water is added to this solution, a black precipitate of silver sulphide 
is formed and the solution then contains dilute nitric acid :_ 

2 NO, 2 -OAg + HoS = Ag 2 S + 2 N 0 2 - 0 H. 

In the same manner the gas produces a yellow precipitate of 
arsenious sulphide in an acid solution of arsenious acid, an orange 
coloured precipitate of antimonous sulphide in a solution of anti- 
monous chloride, &c. 

The affinity of sulphur for most of the metals is so strong and 
for hydrogen so weak that some metals which do not combine 
directly with oxygen can decompose sulphuretted hydrogen. 
Silver, for example, becomes brown in an atmosphere containing 
the compound, owing to the production of silver sulphide. We 
say the silver tarnishes. This action is even produced by the 
minute quantities of the gas exhaled by the human body, contained 
in our coal-gas, and, therefore, always present in inhabited places. 
The darkening of oil paintings by age is also produced by the 
action of sulphuretted hydrogen. The lighter tints in these pic¬ 
tures nearly always contain white-lead (lead carbonate), and this 
becomes gradually converted into black lead sulphide. 

l 2 


148 


Text-Book of Inorganic Chemistry. 


HYDROGEN PERSULPHIDE. 


The composition of this body, which is richer in sulphur than 
sulphuretted hydrogen, has not yet been determined with certainty 
owing to its instability. Its probable composition is H 2 S 2 , corre¬ 
sponding to hydrogen peroxide, H 2 0 2 . 

If hydrochloric acid is added to a solution of potassium penta- 
sulphide, decomposition into potassium chloride, sulphuretted 
hydrogen, and sulphur occurs, as shown on page 145. But if the 
solution of potassium pentasulphide is poured into concentrated 
hydrochloric acid, there is produced, besides the substances 
mentioned above, a heavy oily liquid, with a piercing odour, which 
sinks to the bottom of the acid liquid. If this liquid is removed 
by a separating funnel and left to itself, it gradually decomposes 
into sulphuretted hydrogen and sulphur. 

If the freshly-prepared oil is removed to a thick glass tube and 
then sealed up it undergoes the same decomposition. And the 
greater the quantity of sulphuretted hydrogen set free the greater 
becomes the pressure until finally the gas is liquefied. The sul¬ 
phur which is set free at the same time is usually deposited as 
distinct crystals in the tube. Such tubes containing liquid sulphu¬ 
retted hydrogen are dangerous to handle. It sometimes happens 
that they suddenly explode after having withstood the pressure 
of the gas for years. 


OXYGEN COMPOUNDS OF SULPHUR. 

Sulphur and oxygen unite together in two proportions and 
produce :— 


50 2 . 

5 0 3 . 


Sulphurous anhydride . 
Sulphuric anhydride 


In the former of these sulphur is a tetrad, in the latter it is a 
hexad. Both substances unite with bases and form stable salts, 
and the latter when combined with water yields the most important 
of the acids of sulphur :— 


Sulphuric acid . 




149 


Oxygen Compounds of Sulphur. 

Besides these compounds, which are by far the most important, 
a number of other oxy-acids of sulphur are known, partly in the 
free state and partly combined in their salts. They are all dibasic 
except hydrosulphurous acid. 

Among the next important are :— 

Thiosulphuric acid 

Dithionic acid 

Disulphuric acid 

Thiosulphuric acid is a derivative of sulphuric acid ; in it one 
of the atoms of oxygen united with the hydrogen in sulphuric acid 
is displaced by an atom of dyad sulphur. The two sulphur atoms 
in this compound have, therefore, different functions ; the one 
atom is hexad, and the other, displacing the oxygen, is dyad. 
Dithionic acid is a compound of two atoms of the monad radical 
S 0 2 • OH ; and in disulphuric acid two atoms of the same radical 
are united together by an atom of oxygen. 

Besides these compounds the following are also known, although 
less thoroughly investigated :— 

Hydrosulphurous acid . 

Trithionic acid 

Tetrathionic acid . 

Pentathionic acid . 

In these formulae R stands for an atom of a monad metal in the 
salts of these acids, many of which have not been prepared in the 
free state. 

Hydrosulphurous acid probably contains tetrad sulphur united 
with one atom of oxygen and one of hydrogen to form the monad 
radical (SOH). 

Tri-, tetra-, and pentathionic acids are similarly constituted to 
dithionic and disulphuric acids, and contain two atoms of the 
radical S 0 2 - 0 H united together by one, two, or three atoms of 
dyad sulphur respectively. 


so j 

OR 

H 

cl SO, 

•OR 

s |so* 

•OR 

C /SO,' 

OR 

S MS0 2 ' 

■OR 

c iso. 

•OR 

Mso, 

•OR 


{OH 


S 0 2 

i SO 
(SO 
(sn .oh 


\ SH 

2 -OH 

yOH 



150 


Text-Book of Inorganic Chemistry. 


SULPHUROUS ANHYDRIDE 
Composition : S 0 2 . 

This compound, which is produced when sulphur burns in the 
air or oxygen, is a colourless gas with a powerful piercing odour, 
an acid reaction and taste. Its specific gravity is 2-216, and it 
contains its own volume of oxygen ; for if sulphur is burnt in a 
closed volume of oxygen, the volume of the gas remains unchanged. 
From these data the composition of the gas may be easily calcu¬ 
lated, thus :— 

2 vols. (1 molecule) sulphurous anhydride weigh 2-216 x 28 • 88 = 64 
2 „ „ oxygen „ 2x16 =32 

1 atom of sulphur. 32 

— i.e. a molecule of sulphurous anhydride contains one atom of 
sulphur and two atoms of oxygen. 

Sulphurous anhydride may be easily condensed by cold or 
pressure to a colourless mobile liquid, which boils at — io° and 
solidifies under — 70°. 

The gas is absorbed by water in considerable quantities. One 
volume of water at the ordinary temperature absorbs about fifty 
times its volume of the gas. The aqueous solution possesses the 
odour of the gas and again gives it up when heated. 

Sulphurous anhydride is contained in the gases of volcanoes. 
It is obtained in large quantities for technical purposes by burning 
sulphur or by roasting sulphurous ores (eg. iron pyrites). The 
gas so obtained is of course very impure ; it contains all the nitro¬ 
gen present in the air employed. 

The pure compound is best obtained by reducing sulphuric 
acid with some suitable metal—usually copper. Pieces of copper 
are placed in a large flask, covered with concentrated sulphuric 
acid, and the mixture gently heated until the evolution of gas 
begins (fig. 46). The two substances do not act upon one another 
in the cold. The gas is purified by passing it through a wash- 
bottle containing concentrated sulphuric acid, and is then collected 
in suitable vessels. As it is absorbed in such large quantities by 
water it cannot well be collected over this liquid ; mercury may, 
however, be employed. It can also be easily collected by displace¬ 
ment by leading the gas to the bottom of the vessel to be filled 


Sulphurous Anhydride. I5 1 

and loosely closing the mouth. The heavy gas gradually fills the 
jar and expels the lighter air before it. 

Liquid sulphurous anhydride may be easily obtained in quantity 
by leading the pure gas into a tube surrounded with a freezing 
mixture of ice and salt; the tube (as shown in the figure) being 
contracted at one point so that it may be easily melted and hermeti¬ 
cally sealed after a sufficient quantity of the gas has been con¬ 
densed. 



Fig. 46. 


The production of sulphurous anhydride from sulphuric acid and 
■copper may be expressed by the following equations, which indicate 
that an atom of copper first abstracts an atom of oxygen from the 
sulphuric acid,forming sulphurous acid and copper oxide, the former 
breaking up immediately into sulphurous anhydride and water, and 
the latter uniting with the excess of sulphuric acid to form copper 
sulphate and water :— 

so 2 {8h + Cu - s0 * + H *° + Cu °- 

S°»{oH + Cu0 " SC V°* Cu * H «°' 

The final products are therefore sulphurous anhydride, copper 
sulphate, and water, and the whole reaction may be thus ex¬ 
pressed :— 


















































152 


Text-Book of Inorganic Chemistry. 


2S °,JOH + Cu - SO, + S0 2 • 0 2 Cu + 2 H 2 0 . 

The student will notice that this process is essentially different 
from that which takes place when dilute sulphuric acid acts upon zinc 
to form zinc sulphate and hydrogen. In the latter case the two 
atoms of hydrogen in sulphuric acid are simply displaced by the 
equivalent quantity of zinc (one atom) : — 

S 0 4oH + Zn = S0 2 *0 2 Zn + H 2 ; 

while in the former case the sulphuric acid loses an atom of oxygen 
and becomes reduced to sulphurous acid. 

Only a few metals besides copper can reduce sulphuric acid in 
this manner ; among them are mercury and silver. 

Sulphurous anhydride mixed with carbonic oxide and carbonic 
acid may also be cheaply obtained in large quantities by heating 
concentrated sulphuric acid with charcoal 

so MOH + C = SO 'i + CO + h 3 o. 

This method is valuable for those purposes where the admixture 
with carbonic oxide and acid has no injurious effect—for example, 
in the manufacture of the sulphites. 

Sulphurous anhydride is a compound radical called sulphuryl , 
and behaves in its compounds like a dyad element. Mixed with 
oxygen and led over heated platinum in a finely divided state the 
two substances unite to form sulphuric anhydride :— 

SO >2 + O = S0 3 . 

Its aqueous solution absorbs oxygen from the air and becomes 
converted into sulphuric acid. It further unites directly with 
chlorine, when a mixture of the gases is exposed to bright sunlight, 
producing sulphuryl chloride —S0 2 C1 2 . This compound is a 
colourless liquid with a powerful odour, which boils at 70 ° and is 
decomposed by water into hydrochloric and sulphuric acids. 

From a saturated solution of the gas in water at o°a crystalline 
compound separates out, having the composition SO a + i5H 2 0. 
It melts at 4 0 , and is decomposed into its constituents. 

Most bodies which easily give off oxygen oxidize sulphurous 
anhydride, in the presence of water, to sulphuric acid. It is 
absorbed by manganese peroxide forming manganous sulphate 
and by lead peroxide producing lead sulphate : — 

S0. 2 4 Pb0 2 = S0 2 -0 2 Pb. 


153 


Sulphurous A nhydride. 

The production of this lead sulphate is accompanied with a large 
evolution of light and heat. If a little dried lead peroxide enclosed 
in a piece of thin muslin is brought into a jar of the dry gas, 
the black oxide becomes red-hot and is rapidly converted into 
white lead sulphate. 

Under ordinary circumstances hydrogen has no action on sul¬ 
phurous anhydride, but in the nascent state, especially in the 
presence of acids, the latter is reduced to sulphuretted hydrogen 
and sulphur. If a little of the aqueous solution is poured into a 
flask containing zinc and sulphuric acid, the hydrogen which is 
evolved soon smells strongly of sulphuretted hydrogen and pro¬ 
duces a black precipitate of lead sulphide if led into a solution of lead 
acetate, at the same time the liquid in the flask becomes milky from 
separated sulphur. The sulphur which is here set free is produced 
by the action of sulphuretted hydrogen upon sulphurous anhydride. 
Whenever these two substances come into contact with one 
another, whether in the gaseous state or in solution in water, 
mutual decomposition ensues into sulphur and water :— 

2 H 2 S + SO, = 2H 2 0 + 3 S . 4 

When the reaction takes place slowly the sulphur is often deposited 
in the crystalline form, and it is possible that considerable quan¬ 
tities of the sulphur occurring free in nature have been produced 
by this reaction. 

Sulphurous anhydride bleaches organic colouring matters like 
chlorine, but the bleaching effect is produced in a different way. 
We have seen that chlorine bleaches partly by the formation of 
colourless substitution compounds, and partly by oxidation in the 
presence of water, or by both processes together, and the bleached 
colours cannot be therefore restored. But it appears as if the 
bleaching of sulphurous anhydride were produced by its direct 
union with the colouring substances. And as sulphurous anhydride 
can be expelled from its compounds by strong acids, the colour of 
many bodies bleached with this substance may be restored by 
treatment with strong acids. A rose which has been bleached by 
sulphurous anhydride regains its colour when washed with water 
and placed in dilute sulphuric acid. Similarly, a substance which 
unites with sulphurous anhydride more powerfully than the colour¬ 
ing matter {e.g. a strong base), will also restore the colour. Flannel 
which has been repeatedly washed regains the original yellow 
colour of the wool owing to the action of the alkali contained in the 


154 


Text-Book of Inorganic Chemistry. 

soap used in cleansing it. Substances upon which chlorine acts 
too energetically—silk, wool, straw, &c.—are usually bleached by 
sulphurous anydride. 

Sulphurous anhydride also possesses powerful antiseptic proper¬ 
ties. It destroys the smell of decomposing organic substances and 
stops the action of those organisms which produce fermentation 
and putrefaction. Burning sulphur was used as early as Homer’s 
time to disinfect closed spaces. 1 


Sulphurous acid —probably SO j qj-j ts known only in the 

sulphites. It is a weak dibasic acid, and forms, therefore, two 
series of salts, normal and acid— e.g. :— 


jONa 

Normal sodium sulphite . . . . bu joNa 

, , . crk (OH 

Acid sodium sulphite . . . . ou joNa 


The sulphites are formed by the union of sulphurous anhydride 
with strong bases. They are all decomposed by dilute mineral 
acids, with liberation of sulphurous anhydride. 

The presence of sulphurous acid in any of its compounds may 
be detected by the odour of the gas when liberated by a dilute acid. 
If the quantity is too small to be detected in this way, the gas may 
be allowed to act upon iodic acid, when sulphuric acid and free 
iodine are produced. If a piece of paper moistened with a solution 
of iodic acid and a little starch is hung up in a cylinder containing 
the liquid to be examined, the paper soon becomes blue, owing to 
the union of the liberated iodine with the starch. 

The sulphites may, however, have a different composition to that 
given above. They may be formed on the type of the hypothetical 


(OH 

acid S0 2 -| , which would be called hydrosulphuric acid. This 

acid would then be a monobasic acid, but in which the second 
atom of hydrogen—that united directly to the sulphur—could also 
be displaced by a metal. What we now call acid sodium sulphate 


1 ' Bring sulphur straight and fire ’ (the monarch cries) : 

She hears, and at his word obedient flies. 

With fire and sulphur, cure of noxious fumes, 

He purg’d the walls and blood-polluted rooms. 

Odyssey, xxii. 527-530. (Pope’s Translation.') 


Sulphuric Acid. 


155 


would then be the normal hydrosulphate : S0 2 j ^^ a , and our normal 

sulphite would be sodium sodiosulphate : S0 2 j a . These con¬ 
siderations are supported by many facts in organic chemistry, 
and the decision which view is correct cannot be long delayed. 
Possibly isomeric salts of the two acids exist together. 


Corresponding to this hypothetical hydrosulphuric acid with 
one atom less oxygen is :— 

(OH 

Hydrosulphurous acid : SO j ^ , sometimes called after its dis¬ 
coverer, Schiitzenberger’s Acid. The acid, which has not yet been 
obtained in the pure state, is prepared by digesting a concentrated 
aqueous solution of sulphurous anhydride with metallic zinc, in 
which the latter dissolves with the evolution of gas. The yellow- 
liquid so obtained, which contains hydrosulphurous acid, possesses 
powerful reducing and bleaching properties. It rapidly decomposes 
first into thiosulphuric acid and then into sulphurous anhydride, 
water, and sulphur. The sodium salt may be prepared by digest¬ 
ing together zinc and a concentrated solution of acid sodium 
sulphite. 


SULPHURIC ACID. 

(OH 

Composition : H 2 S0 4 , or S0 2 j 

If a further atom of oxygen is made to unite with sulphurous 
anhydride, the compound sulphuric anhydride (S0 3 ) is formed. 
But far more important than this body is the substance produced 
when it unites with water—viz. sulphuric acid, or oil of vitriol, it 
is sometimes called from the old method of preparing it. 

Sulphuric acid is a compound of hexad sulphur, and contains the 
radical S0 2 united to two atoms of hydrogen through the interven¬ 
tion of two oxygen atoms. 

Sulphuric acid is a colourless, odourless, viscid liquid, which 
does not fume in the air, and has a specific gravity of 1 * 84 . It boils 




156 Text-Book of Inorganic Chemistry. 

at 338 ° and can be distilled, but only with partial decomposition, into 
sulphuric anhydride and water. It mixes with water in all propor¬ 
tions, and during mixing evolves a large quantity of heat. Even 
when very largely diluted with water it reacts and tastes strongly 
acid. Below o° it freezes to a crystalline solid. 

In nature sulphuric acid occurs chiefly in combination with lime 
as gypsum or anhydrite , widely distributed and in very large quanti¬ 
ties, and besides this as heavy spar or barite (barium sulphate), as 
celestine (strontium sulphate), as Glauber's salt (sodium sulphate), 
&c. 

Calcium sulphate is such a common mineral that it might be 
thought capable of yielding an inexhaustible source of cheap 
sulphuric acid, and this would be so were it possible to separate 
the sulphuric acid from its combination with the lime as easily as 
we can separate carbonic acid from chalk or nitric acid from nitre. 
But sulphuric acid is so firmly united with the lime in calcium 
sulphate that we can neither expel it by heating nor by the action of 
a stronger acid, and gypsum is not therefore adapted for the manu¬ 
facture of sulphuric acid, nor has it ever been used for this purpose. 

The acid may, however, be obtained in many other ways. 
Sulphur, which burns in oxygen to sulphurous anhydride, is con¬ 
verted when heated with concentrated nitric acid or other powerful 
oxidizing agents into the higher oxide—sulphuric acid. And here 
the question may present itself, Why is it that while carbon and phos¬ 
phorus when burning in the air always produce their higher oxide, 
sulphur only forms its lower oxide ? The reason is that sulphuric 
anhydride cannot exist at a high temperature, but is decomposed 
into oxygen and its lower oxide, and that the temperature of com¬ 
bustion of sulphur is higher than the point at which sulphuric anhy¬ 
dride is decomposed. If we could by any means lower the tempe¬ 
rature of combustion of sulphur, it would undoubtedly produce 
sulphuric anhydride when burning in air, but up to the present 
this problem has not been solved, and we must therefore look for 
some other method of oxidation. The best substance for this 
purpose, and that exclusively used for the manufacture of sulphuric 
acid, is nitric acid. 

In order to understand this process, it must be remembered 
that nitric acid, which has the composition N0 2 -OH, yields up 
a portion of its oxygen and becomes converted into nitric 
oxide (NO) in the presence of sulphurous anhydride and water. 
Three molecules of sulphurous anhydride therefore require two 


Sulphuric Acid. 157 

molecules of nitric acid and two of water to be converted into sul¬ 
phuric acid:— 


3 S0 2 


2 N 0 „- 0 H 


2 H s O = 3S0 2 |°g + 2NO. 


But if as in this equation two mojecules of nitric acid were con¬ 
sumed in the production of every three molecules of sulphuric acid, 
the price of the latter acid would be scarcely lower than that of 
the former. As a matter of fact, the nitric acid can oxidize much 
larger quantities of sulphurous acid than that expressed in the 
above equation—more than ten times as much. This is rendered 
possible by the remarkable property which the colourless nitric 
oxide possesses of uniting with the oxygen of the air, even at the 
ordinary temperature, and forming red vapours consisting of nitrous 
anhydride and nitric peroxide, and these gases again give up oxygen 
to sulphurous anhydride, oxidizing it to sulphuric acid in the pre¬ 
sence of water, and being themselves again reduced to nitric oxide. 
It thus appears as if a small quantity of nitric oxide would be able 
to convert an unlimited quantity of sulphurous anhydride into sul¬ 
phuric acid, if allowed to act upon a mixture of the gas and pure 
oxygen together with water vapour in the correct proportions. The 
oxygen would at once convert the nitric oxide (NO) into nitric per¬ 
oxide (NO,), which would be again reduced to nitric oxide in the 
presence of sulphurous anhydride and water vapour . 

NO, + S0 2 + H s O = S0 2 |q^ + NO, 

and this could then produce further quantities of sulphuric acid, 
and so on. 

Blit such a process cannot be carried on in practice, simply 
becauses it presupposes an inexhaustible supply of pure oxygen. 
The manufacturer must make use of the oxygen contained m the 
air. The air, by the gradual abstraction of its oxygen, becomes 
largely diluted with nitrogen, and ultimately the small proportion 
of oxygen present will not unite with the nitric oxide sufficiently 
ouickly. When this point is reached, the valuable nitrogen com¬ 
pounds would be'wasted if a method had not been discovered of 
condensing them, and reintroducing them into the chambers. 

The manufacture of sulphuric acid on a large scale, based 
upon the above-mentioned chemical principles, is conducted as 

“'The sulphurous anhydride is obtained either by burning sulphur 


15 S Text-Book of Inorganic Chemistry. 

or by roasting iron pyrites. It is then allowed to stream, mixed 
with atmospheric air, into a large Leaden chamber, where it comes 
into contact with nitric acid vapours. A small portion of the sul¬ 
phurous anhydride is at once oxidized by the nitric acid to sulphuric 
acid, but by far the greater portion passes on, mixed with nitric 
peroxide, nitrous anhydride, and nitric oxide, into the next chamber, 
into which steam is led from a boiler. The reactions mentioned 
above go on in this chamber, and the dilute sulphuric acid collects 
on the floor. 1 

The motion of the gases through the chambers is produced by 
a tall shaft placed at the further end, which returns the useless 
nitrogen to the atmosphere. But before it reaches the shaft it is 
made to pass through what is called a Gay Lussac’s tower to retain 
the valuable oxides of nitrogen so that they may be again used to 
produce a further quantity of sulphuric acid. It has been found 
that strong sulphuric acid can absorb nitric peroxide, and especially 
nitrous anhydride, in considerable quantities. This property is 
utilized in the Gay-Lussac’s tower by causing strong sulphuric acid 
to trickle down in a finely divided state over pieces of coke, in 
order to expose as large a surface of the acid as possible to the 
ascending gases. The strong acid containing the oxides of nitro¬ 
gen in solution (so-called nitrated acid), which collects at the foot 
of the tower, is pumped back to the other end of the chambers. 
The gases which it contains are then again liberated by mixing it 
with dilute acid, and are again introduced into the chambers by 
allowing it to flow down a second tower (Glovers tower), through 
which the sulphurous anhydride and air are made to pass. 

The chambers in which the sulphuric acid is produced are 
made of lead, because lead is the only substance available which 
resists to some extent the action of sulphuric acid. The leaden 
plates constituting the chambers are melted together by means of 
the oxy-hydrogen blowpipe, and are supported externally by a stout 
wooden framework. 

The acid which collects on the floor of the chambers—the 
chamber acid— is somewhat dilute, and only contains little more 

1 The reactions given here only explain in a general manner what goes on in 
the chambers. What really happens is doubtful, and is probably dependent 
upon the proportions in which the various substances are introduced into the 
chambers. If the supply of steam is deficient, white crystals having the com¬ 
position S 0 2 | are produced, which under the action of water at once 
break up into sulphuric acid, nitric acid, and nitric oxide.— Ed. 


? 59 


Sulphuric Acid. 

than 64 per cent, of the pure acid (sp. gr. = 1*55). By heating in 
leaden vessels it is concentrated until - it contains about 78 per cent, 
of the pure acid. This is the brown acid of commerce, and has a 
specific gravity of about 171. The concentration cannot be carried 
further in leaden vessels, as at this point the lead begins to be 
attacked. Further concentration is. then carried on in vessels of 
glass or better of platinum. The sulphuric acid so prepared is 
usually coloured brown from' traces of organic substances, and 
always contains lead sulphate in solution. If diluted with water or 
alcohol, the liquid becomes turbid, and ultimately deposits a white 
precipitate of this lead sulphate, which is more soluble in the con¬ 
centrated than in the dilute acid. 

If iron pyrites is used in the manufacture of sulphuric acid, 
the acid always contains arsenic, sometimes in not inconsiderable 
quantities. A sample of the acid containing arsenic when diluted 
with water, and saturated with sulphuretted hydrogen, first turns 
yellow, and then deposits a yellow precipitate of arsenious sulphide. 
By far the greater quantity of sulphuric acid which is brought into 
trade contains arsenic ; only a comparatively small quantity of the 
more expensive acid free from arsenic is manufactured from Sici¬ 
lian sulphur. 

The colourless, concentrated, chemically pure sulphuric acid 
is prepared by distilling the acid free from arsenic in platinum 
vessels. During distillation it always undergoes a partial decom¬ 
position into sulphuric anhydride, which passes over with the 
distillate and water which remains behind ; this goes on until the 
acid contains 98 per cent, of the pure compound, which then 
distils over unchanged. 

Sulphuric acid is one of the strongest acids, and expels there¬ 
fore nearly eveiy other acid from its compounds. It is, like 
sulphurous acid, a dibasic acid, and its acid salts possess a strong 
acid reaction and taste. With few exceptions, its salts are soluble 
in water. Sulphuric acid is distinguished by its powerful attraction 
for water, with which it unites in several definite proportions, 
always evolving a large quantity of heat. The compound with one 

molecule of water—S0 2 |q^ + H 2 0 has the specific gravity 

178, and solidifies in a crystalline form at + 8°. Unlike water, it 
does not expand on solidification, and vessels filled with this acid 
do not therefore burst on freezing. If to this compound a further 


160 Text-Book of Inorganic Chemistry . 

quantity of water is added, a further quantity of heat, though less 
than before, is evolved. And the quantities of heat which are set 
free when one, two, or three molecules of water unite with one 
molecule of sulphuric acid are in a certain definite proportion to 
one another. 

The strong attraction of sulphuric acid for water is employed 
to dry those gases upon which the acid has no action (eg. oxygen, 
hydrogen, carbonic acid), the gases being simply led through a 
wash-bottle containing the strong acid. It also serves to dry solid 
and liquid bodies which are placed in a closed space (a desiccator) 
with the concentrated acid. The acid then rapidly absorbs all 
the water vapour which evaporates from the substances to be dried. 

Ordinary commercial sulphuric acid is always of a brownish 
colour, due to the decomposition of particles of dust of organic 
origin. If a splinter of wood is dipped into the concentrated acid 
it soon becomes brown and then black from the charcoal which 
the acid has set free from the wood by abstracting oxygen and 
hydrogen in the form of water. These elements are not present 
as water in the wood, but the acid compels them to combine to 
form water, with which it then unites. 

Sulphuric acid is by no means so permanent as its powerful 
affinities and high boiling point might lead us to suppose. Even 
below a red heat, if allowed to drop upon hot bricks or upon pieces 
of pumice stone contained in a hot platinum retort, it is decom¬ 
posed into water, oxygen, and sulphurous anhydride. If the 
mixture of. gases so obtained is led into a solution of caustic soda, 
the sulphurous anhydride is absorbed, and considerable quantities 
of pure oxygen may be obtained. The instability of sulphuric 
acid at high temperatures is the reason why burning sulphur 
produces sulphurous anhydride and not sulphuric anhydride. 

Nascent hydrogen reduces concentrated sulphuric acid (not the 
dilute acid) to sulphuretted hydrogen :— 

S ° 4 oH + 4H * = H * S + 4 H 2°- 

If a few drops of concentrated sulphuric acid are allowed to 
flow by means of a funnel tube into a flask evolving hydrogen, the 
hydrogen soon acquires the characteristic odour of sulphuretted 
hydrogen, and produces a black precipitate of lead sulphide in a 
solution of lead acetate. Zinc, iron, nickel, manganese, and other 
metals which possess a strong attraction for oxygen and are 


Sulphuric Acid. 


161 


dissolved by dilute sulphuric acid with evolution of hydrogen, are 
not attacked even when heated with the concentrated acid, probably 
because their sulphates are so insoluble in the concentrated acid 
that they incrust the metals and prevent any further action. 
Other metals, as copper and mercury, which in consequence of a 
feebler attraction for oxygen are not attacked by dilute sulphuric 
acid, reduce the concentrated acid when heated with it, forming 
sulphurous anhydride and a sulphate soluble in concentrated 
sulphuric acid. In this way sulphuric acid, which we prepare 
from impure sulphurous anhydride, affords us the best material for 
the preparation of this substance in the pure state. 

Sulphuric acid contains two atoms of displaceable hydrogen, 
and is therefore a dibasic acid. With monad metals it forms two 
series of salts—the normal and acid— e.g. :— 



Normal sodium sulphate . 
Acid sodium sulphate 


Most of the sulphates are soluble in water, the most important 
exception being barium sulphate, which is quite insoluble in water 
and dilute acids, and which is therefore used to detect sulphuric acid 
or a sulphate in an aqueous solution. On the addition of a few drops 
of barium chloride to such a solution, the presence of sulphuric 
acid is at once recognized by a white turbidity or precipitate of 
barium sulphate which is insoluble in hydrochloric acid. 

The great progress which has been made during the past fifty 
years in technical chemistry and the arts generally, is due to a 
large extent to improvements in the manufacture of sulphuric acid. 
Hundreds of thousands of tons of sulphuric acid are manufactured 
annually in England, Germany, and France alone, and by far the 
greatest part of this immense quantity is used for the production of 
soda, from which again two of the most indispensable articles 
of daily life—soap and glass—are obtained. Besides this the 
acid is also used for an immense number of other chemical pro¬ 
cesses, and, in fact, there is scarcely any chemical manufacture in 
which sulphuric acid does not take a direct or indirect part. It 
will, therefore, be at once apparent that a reduction in the price of 
this important substance to one-tenth the amount paid for it a little 
more than a century ago must have exercised a most beneficial in¬ 
fluence on the development of technical chemistry. 

If, as is not only possible but even probable, further improve- 


M 


152 Text-Book of Inorganic Chemistry. 

ments are made in the manufacture of sulphuric acid, and its price 
becomes still lower, the price of soda, glass, soap, and stearine, o 
superphosphate manures and hundreds of other things, would at 
once fall to a considerable extent. 


SULPHURIC ANHYDRIDE. 

Composition : SO s . 

This substance, which is sometimes called sulphuric acid, is very 
unlike sulphuric acid in its properties. It is a solid body, crystal¬ 
lizing in white, silky needles. It melts at 15 0 and boils at 46 , 
producing vapours which form thick white fumes of sulphuric acid 
in moist air. The density of its vapour, compared with the air as 
unity, is 277, corresponding to a molecular weight of 80. 

Sulphuric anhydride may be easily obtained from sulphurous 
anhydride and oxygen. A mixture of the two gases remains un¬ 
changed even if heated or exposed to sunlight, but if led over a 
layer of heated platinum, contained in a tube of hard glass, the two 
substances at once unite with one another, and the presence of sul¬ 
phuric anhydride is made manifest by the dense white fumes pro¬ 
duced where the gases come into contact with the air. The 
platinum itself remains quite unchanged ; it acts in some unknown 
manner upon the gaseous mixture, and its action is an example of 
what is called, for want of a better name, contact action. Perhaps 
the phenomenon depends upon the property of platinum to con¬ 
dense considerable quantities of oxygen on its surface and so liquefy 
the gas ; or it may be that the oxygen is converted into its active 
modification—ozone. Instead of pure platinum it is better to em¬ 
ploy platinized asbestos— i.e. asbestos of which the surface has been 
covered with a thin layer of platinum. 1 

Sulphuric anhydride is commonly prepared from Nordhausen 
or fuming sulphuric acid, which may be considered as a solution of 
the anhydride in sulphuric acid. When this acid is gently heated 
in a retort, the anhydride volatilizes and condenses in a cool dry 
receiver as a colourless crystalline mass. 

Sulphuric anhydride is especially characterized by its strong 
attraction for water. If a drop of the molten substance is allowed 

1 Prepared by dipping asbestos into platinic chloride solution, drying, and 
heating to redness.—E d. 



Nordhausen or Fuming Sulphuric Acid. 163 

to fall into a vessel of water, it sets free a large quantity of heat and 
hisses like a red-hot iron on immersion in water. If a drop of 
water happens to fall into a glass vessel containing the anhydride 
chemical union at once takes place with an explosion, and the 
vessel is always shattered. In all cases the compound produced 
by its union with water is sulphuric acid. Sulphuric anhydride 
acts even more powerfully on organic substances than sulphuric 
acid ; it abstracts water and chars them. A cork or india-rubber 
stopper cannot, therefore, be used to close a bottle containing 
the volatile anhydride. Formerly sulphuric anhydride was only 
employed for certain reactions in the laboratory on a small scale, 
but it is now used more and more in the arts; for example, in the 
manufacture of alizarine from anthracene. It is now prepared in 
considerable quantities in chemical works, and is an article of 
commerce. 


NORDHAUSEN OR FUMING SULPHURIC ACID. 

This acid is not a definite chemical compound, but consists (as 
we have previously stated) of sulphuric acid containing more or less 
sulphuric anhydride in solution. In former times it was principally 
manufactured at Nordhausen, in the Harz Mountains, though at 
present none is made there. 

Fuming sulphuric acid is a thick oily liquid, usually coloured 
brown from the presence of minute particles of carbon, and is dis¬ 
tinguished from ordinary sulphuric acid by the fact that it fumes 
in the air. These fumes consist of the volatile sulphuric anhydride 
which it gives off, and which unites with the moisture of the air. 

The acid is usually prepared from green vitriol (ferrous sul¬ 
phate). This salt crystallizes with seven molecules of water, and 
has the composition: S 0 2 - 0 2 Fe + 7 H 2 0 . When heated in the 
air it easily loses six molecules of water, but the seventh only at a 
higher temperature, at which the salt itself begins to be decomposed. 
The salt, after being dried as thoroughly as possible, is heated in 
clay retorts provided with receivers of the same material. It is 
then decomposed into ferric oxide (called caput mortuum or col- 
cothar ), which remains behind in the retorts, and sulphurous anhy¬ 
dride with sulphuric anhydride, which distil over, the latter con¬ 
densing in the cool receivers. The small quantity of water which 
the dried salt still contains suffices to convert the sulphuric anhy- 

M 2 



1 64 Text-Book of Inorganic Chemistry. . 

dride in the retort into liquid fuming sulphuric acid. The following 
equation represents the process :— 

2S0 2 -0 2 Fe = Fe 2 0 3 + S 0 2 + S 0 3 . 

A larger quantity of acid is obtained, and without the production 
of sulphurous anhydride, when the mother-liquor from ferrous sul¬ 
phate, which contains the ferric salt, is evaporated and calcined. 
The mass so obtained consists principally of basic ferric sulphate 
of the composition: S 2 O p Fe 2 = 2SO s , Fe 2 0 3 , which on further heat¬ 
ing in retorts breaks up into ferric oxide and sulphuric anhydride. 

As was stated on p. 162, fuming sulphuric acid when heated in 
a retort boils and gives off the sulphuric anhydride which it con¬ 
tains even below ioo°. As the quantity of anhydride becomes less 
and less the boiling point rises, until at last ordinary sulphuric acid 
remains in the retort. 

When cooled, fuming sulphuric acid deposits colourless crystals, 
which melt at 35 0 , and consist of disulphuric acid, the properties 
of which are described below (p. 166). 

Sulphurous anhydride can unite not only with an atom of 
oxygen but also with an equivalent quantity— i.e. two atoms—of 
chlorine. The compound so produced is Sulphuryl chloride : 
S 0 2 C 1 2 . Besides this body another chlorine compound is known, 
intermediate between sulphuryl chloride and sulphuric acid, and 
which may be considered as sulphuric acid with one atom of 
hydroxyl (OH) displaced by one atom of chlorine. This is Chlor- 

sulphonic acid : S 0 2 j q R . The latter is a monobasic acid, because 

it contains one atom of hydrogen displaceable by a metal, while 
the former is an indifferent substance, and cannot enter into com¬ 
bination with bases. 


SULPHURYL CHLORIDE. 

Composition : S 0 2 C 1 2 . 

This substance is a mobile colourless liquid, with a piercing 
odour. It fumes slightly in the air, has a specific gravity of 17, 
and boils at 70°. It may be obtained by direct union of chlorine 
and sulphurous anhydride when a mixture of the two gases in 
equal volumes is exposed to direct sunlight, or more readily, and 
in larger quantities, when chlorsulphonic acid is heated in closed 



Sulphuryl Chloride. 


165 


tubes to 200° or 210° for about twelve hours. Two molecules of 
chlorsulphonic acid are then decomposed into one molecule of 
sulphuryl chloride and one of sulphuric acid :— 



The liquid product, which is usually of a greenish yellow colour, 
is heated in a retort, and the portion passing over up to no° again 
rectified on the water bath. The portion which then distils over 
at about 70° is nearly pure sulphuryl chloride. In order to separate 
minute quantities of chlorsulphonic acid and sulphuric anhydride, 
which it still contains, the distillate is poured into a separating 
funnel containing pieces of ice. The ice at once acts upon these 
impurities, while the sulphuryl chloride, which is scarcely attacked, 
sinks as a heavy oil to the bottom of the funnel. It is then freed 
from water with phosphoric anhydride, redistilled, and if the first 
portions of the distillate, which contain sulphurous acid and 
chlorine, are rejected, is so obtained perfectly pure. 

Sulphuryl chloride is easily decomposed by water into sulphuric 
and hydrochloric acids :— 



CHLORSULPHONIC ACID. 


' Cl 

Composition : S 0 2 j 


This substance is also a colourless liquid, with a specific gravity 
of 177, and boiling at 153 0 . It is at once decomposed by water, 
with a considerable evolution of heat, into sulphuric and hydro¬ 
chloric acids :— 



+ HC1. 


Chlorsulphonic acid may be produced by the direct union of 
sulphuric anhydride and hydrochloric acid gas—S 0 2 0 + HC 1 = 

S 0 2 j qjj, but is best obtained by heating together molecular 

weights of sulphuric acid and phosphorus pentachloride, when 
hydrochloric acid and phosphoric oxychloride are also produced 



+ POCI 3 + HC1. 



166 Text-Book of Inorganic Chemistry. 

It may then be separated from the more volatile phosphoric oxy¬ 
chloride (B.P, = no°) by fractional distillation. Metallic salts of 
chlorsulphonic acid have not yet been prepared. 


NITROSULPHONIC ACID. 

Composition : S 0 2 

This compound, which may be considered as sulphuric acid in 
which one atom of hydroxyl is displaced by the group of atoms 
N0 2 , separates out as colourless crystals when nitric peroxide is 
led into concentrated sulphuric acid or sulphurous anhydride into 
cold fuming nitric acid. The latter reaction is expressed by the 
equation 

S 0 2 + NO s OH = S 0 2 {g°* 

Nitrosulphonic acid is also produced by the action of sulphurous 
anhydride on a mixture of nitric oxide and oxygen in the presence 
of a small quantity of water, and occurs therefore in the sulphuric 
acid chambers when the quantity of steam is insufficient. The 
so-called white crystals to which we have previously referred 
(p. 158 note), consist of this compound, and their presence in the 
chambers is usually considered to indicate irregularities in the 
manufacture. 

Nitrosulphonic acid is dissolved by sulphuric acid, but decom¬ 
poses when gently heated or when brought into contact with water. 
In the latter case sulphuric acid and nitrous acid are produced, and 
the latter then decomposes into nitric oxide and nitric acid. When 
dissolved in sulphuric acid it requires a greater quantity of water 
to decompose it than when in the free state. 


DISULPHURIC ACID. 

Composition : H 2 S 2 0 7 = O | sq 2 q[] 

The latter formula indicates that two atoms of the sulphuric 
acid radical (sulphuryl = S 0 2 ) are contained in a molecule of di- 
sulphuric acid, that they are united together by an atom of oxygen, 




Disulphuric A cid. 


167 


and that the other two unsatisfied bonds of the S 0 2 are connected 
with two atoms of hydroxyl (HO), the hydrogen of which can be 
displaced by metals ; the acid is therefore dibasic. 

Disulphuric acid is deposited from fuming sulphuric acid at a 
low temperature as colourless and often large-sized crystals, which 
melt at 35 0 and fume in the air. it is decomposed when brought 
into contact with water into ordinary sulphuric acid :— 


O 


so 2 -oh 

so 2 -oh 


+ h 2 0 


2 S 0 2 


OH 

OH 


and when heated into sulphuric anhydride and sulphuric acid :— 


O 


(so 2 -oh 

(S 0 2 - 0 H 


S 0 3 + 


SO 


jOH 
2 (OH 


Its salts are more stable than the free acid. The sodium salt is 
easily obtained by heating acid sodium sulphate as long as water 
is expelled. The following equation shows the reaction 


2 S 0 2 


ONa 

OH 


Acid sodium 
sulphate 


n fS 0 2 - ONa 
u (S 0 2 -ONa 

Sodium 

disulphate 


+ 


h 2 o. 


This salt dissolves in water unchanged. When strongly heated it 
breaks up into normal sodium sulphate and sulphuric anhydride 


n j S 0 2 -ONa 
U (S 0 2 . 0 Na 


= SO, 


ONa 

ONa 


SO, 


THIOSULPHURIC ACID. 

Thiosulphuric acid is not known in the free state, but only in 
its salts. It is a dibasic acid, and is to be considered as sulphuric 
acid, in which one of the atoms of oxygen united with the hydrogen 
is displaced by dyad sulphur. The composition of the hypothetical 

acid is therefore SO, j or H 2 S 2 0 3 . Before the relations of thio- 

sulphuric acid to sulphuric acid were known, the acid was compared 
with dithionic acid, which has the empirical formula, H 2 S 2 0 6 . 
Whence arose the former name for this acid : hyposulphurous 
acid , x based upon the old name for dithionic acid—hyposulphuric 

acid. 

1 This name survives in the commercial designation for the sodium salt • 
hyposulphite of soda.—E d. 



168 


Text-Book of Inorganic Chemistry. 


Of the compounds of thiosulphuric acid we are only acquainted 
with its metallic salts, some of which are stable substances, and of 
which the sodium salt is most easily obtained in the crystalline 
state. If we endeavour to set free the thiosulphuric acid from its 
salts by means of a stronger acid, we find that the acid can only 
exist for a very short time even in dilute solutions. The liquid 
soon becomes turbid from deposited sulphur, and begins to Smell 
of sulphurous anhydride. The acid decomposes in fact into 
sulphur, sulphurous anhydride, and water :— 

SO *{oH = S 0 2 + S + H 2 G. 


This reaction may be used for the detection of the thiosulphates, 
especially as the separated sulphur always possesses a yellow 
colour. The best known salt is sodium thiosulphate, which with 
the potassium salt is distinguished by its stability. This impor¬ 
tant salt may be obtained in various ways. It is produced if 
an aqueous solution of neutral sodium sulphite is boiled with 
sulphur :— 


sol ONa + s 

bU (ONa + b 


so 2 


ONa 

SNa 


or if sulphurous anhydride is led into a solution of sodium mono¬ 
sulphide. Normal sodium sulphite and sulphuretted hydrogen are 
then first produced :— 


SO, + Na,S + H 2 0 . so | ONa + H * S > 

and the latter gas reacts at once on the excess of sulphurous anhy¬ 
dride, forming water and free sulphur :— 

S 0 2 + 2H 2 S = 2 h 2 o + 3 S. 

The greater part of this sulphur sinks to the bottom of the vessel, 
but a portion unites with the sodium sulphite to form sodium thio¬ 
sulphate. 

Sulphur therefore acts on the sulphites in the same way as 
oxygen. Just as a solution of sodium sulphite when exposed to the 
air absorbs oxygen and becomes converted into sodium sulphate, 
so the same solution when warmed with powdered sulphur unites 
with it to form the corresponding thiosulphate. The composition 
of thiosulphuric acid is also interesting because its molecule con¬ 
tains two atoms of sulphur of different valency. The atom in the 
radical sulphuryl (S 0 2 ) is hexad, while that uniting the atom of 
hydrogen or monad metal to this radical is dyad. Sodium thiosul- 


Thiosulphuric Acid. 169 

phate, together with polysulphides, is produced when a solution of 
caustic soda is boiled with an excess of sulphur :— 

6 NaOH + 12S = SO,|ONa + 2 Na a S s + 3 H a O. 

Sodium 

. pentasulphide 

Finally, some sulphides (e.g. that of calcium) are oxidized when 
exposed to the air and converted into thiosulphates. Calcium 
sulphide is largely produced as a bye-product in the manufacture of 
soda, and this alkali-waste when exposed to the air becomes partly 
converted into calcium thiosulphate, which, when digested with a 
solution of sodium carbonate or sulphate, yields insoluble calcium 
carbonate or sulphate, and a solution of sodium thiosulphate. 

The compounds of thiosulphuric acid with the heavy metals 
are mostly insoluble and easily decomposed. Sodium thiosulphate 
produces in a solution of silver nitrate a yellowish white precipitate 
of silver thiosulphate. This salt, however, rapidly darkens in 
colour, and at last becomes quite black, owing to the formation of 
silver sulphide and sulphuric acid :— 

S0 4sa| + H ’° - A & S + S ° 2 {oH 
In the same way the white lead salt blackens on boiling, and for 
the same reason. 

The extraordinary property which sodium thiosulphate, in com¬ 
mon with other soluble thiosulphates, possesses of dissolving chlo¬ 
ride, bromide, and iodide of silver, with which it forms soluble 
double salts, is of considerable importance. Use is made of this 
property chiefly in the production of photographs. After the sen¬ 
sitive plate, which is coated with a thin layer of these compounds, 
has been exposed in the camera, and after the iodide, &c., of silver 
has been decomposed in those parts on which the light has fallen, 
it becomes necessary in order to make the picture permanent to 
remove the iodide, &c., which has remained unchanged. This is 
effected by dipping the plate into a solution of sodium thiosulphate, 
which at once removes the unchanged compounds. Many tons of 
sodium thiosulphate are annually manufactured for this purpose. 



170 


Text-Book of Inorganic Chemistry. 


DITHIONIC ACID. 


This acid, formerly called hyposuiphnric acid , is not known to 
us in the iree state, but only in its salts. From the composition of 
these bodies we consider the acid to be :— 


h 2 s 2 0 6 


f SO„OH 
{S 0 2 0 H 


the latter formula expressing that the molecule of the acid consists 
of two atoms of the radical : ~S 0 2 0 H, each of which contains an 
atom of displaceable hydrogen, thus making the acid dibasic. 

Manganous dithionate may be obtained by leading sulphurous 
anhydride into water containing finely divided manganese peroxide 
in suspension. The dry peroxide unites at once with sulphurous 
anhydride to form manganous sulphate :— 

S 0 2 + Mn 0 2 = S 0 2 - 0 2 Mn, 


but in the presence of water twice as much of the gas takes part 
in the reaction, and manganous dithionate is formed :— 

2 S0 2 + Mn 0 2 = | so 2 'o Mn * 

At the same time a small quantity of the sulphate is always 
produced, and both salts remain in solution. The filtered liquid is 
then mixed with a slight excess of baryta water— i.e. until it be¬ 
comes faintly alkaline—which decomposes the manganous sul¬ 
phate, and precipitates manganous hydrate and barium sulphate. 
The liquid, which now contains barium dithionate and the excess 
of barium hydrate, is again filtered, and the slight excess of the 
latter converted into insoluble barium carbonate by passing a 
stream of carbonic acid, and at the same time gently warming 
it. From the clear solution, barium dithionate is deposited in 
large, colourless crystals on evaporating down, and allowing to 
cool. If lime water is employed instead of baryta water, calcium 
dithionate is obtained instead of the barium salt. By decomposing 
these compounds with the soluble sulphates of other metals, the 
dithionates of these metals may be easily obtained. All the salts 
of dithionic acid are soluble in water. 

The acid itself may be set free from barium dithionate by 
exactly precipitating the barium with sulphuric acid, or from the 
lead salt by precipitating the lead with sulphuretted hydrogen, and 
then filtering off the clear aqueous solution. The water may be 


Dithionic Acid. 


171 


removed to some extent by evaporation in a vacuum over sul¬ 
phuric acid; but if the concentration is driven too far, the acid 
decomposes into sulphuric acid and sulphurous anhydride :— 


(S 0 2 - 0 H 

|so 2 -oh 



The same decomposition, which may be recognized by the odour 
of sulphurous anhydride and by the precipitate produced on the 
addition of a drop of barium chloride, takes place when the 
dilute aqueous solution of the acid is boiled. But although easily 
decomposed, dithionic acid is decidedly more stable than thio- 
sulphuric acid. 

Dithionic acid is usually thought to be a dibasic acid, although 
this has not yet been definitely proved. At present only neutral 
salts of the acid are known, and until we can prepare the acid 
salts which dithionic acid in common with all polybasic acids 
ought to yield, the question of its basicity must remain unsettled. 
Perhaps these compounds are as easily decomposed as the free 
acid, and break up on evaporation into the acid sulphate and 
sulphurous anhydride. It is therefore always possible that the 
molecule of dithionic acid is only half as large as that generally 
accepted, that it is a monobasic acid and a compound of pentad 
sulphur (S 0 2 OH). This possibility is not contradicted by the 
production and chemical behaviour of the acid, nor is it impossible 
to imagine that sulphur might exist in some compounds as a pentad 
element. 


POLYTHIONIC ACIDS. 


Under this name may be included the three acids of sulphur 
which contain more than two atoms of sulphur. These are 
Trithionic Acid : H 2 S 3 O e , Tetrathionic Acid : H 2 S 4 O e , and Penta- 
thionic Acid : H 2 S 3 0 6 . In these acids it is thought that the two 
groups of atoms which constitute dithionic acid are united together 
by one, two, and three atoms of sulphur respectively, as is expressed 
in the following formulae :— 



Tetrathionic acid = 



17 2 Text-Book of Inorganic Chemistry . 


Pentathionic acid = H 2 S 3 0 6 


c ( S0 2 0H 
~ ^3{S0 2 0H 


formulas which involve the supposition that the double as well as 
the triple atom of sulphur can play the part of a dyad radical just 
like a simple atom of the element. 


Trithionic Acid : HoSgOg 


Q (SO,OH 
b (S0 2 OH 


The potassium salt of this acid may be obtained by leading 
sulphurous anhydride into a solution of potassium thiosulphate, 
sulphur being at the same time separated :— 

2 S0 4sk + 3SO * - 2 K A°« + S - 


The acid, which may be set free from the potassium salt by the 
action of fluosilicic acid, very easily decomposes into sulphur, sul¬ 
phurous anhydride, and sulphuric acid. 


Tetrathionic Acid : H 2 S 4 O e = S 3 


SO q OH 

SO'OH 


The sodium salt is obtained by adding iodine to an aqueous 
solution of sodium thiosulphate :— 

2S ° ! jsNa + h - Na 2 S 4 0 6 + 2 NaI. 

The barium salt may also be prepared in the same way. If this 
compound is decomposed with sulphuric acid, free tetrathionic 
acid is produced, which decomposes at a certain stage of concen¬ 
tration into sulphur, sulphurous anhydride, and sulphuric acid. 
Its salts, which are mostly soluble in water, are decomposed on 
evaporation into sulphur and the corresponding trithionates. 


Pentathionic Acid 1 : H 2 S 5 0 6 = Ssjcj^OH 

The free acid is produced together with large quantities of sulphur 
when sulphuretted hydrogen is led into an aqueous solution of 
sulphurous acid :— 

5 H 2 S + 5S0 2 = H 2 S 3 0 6 + 5 S + 4H 2 0. 

The barium salt may also be prepared by adding sulphur di- 

1 Recent experiments have shown that the existence of pentathionic acid is 
doubtful. 


Compounds of Sulphur and Chlorine . 173 

chloride to water containing finely divided barium thiosulphate in 
suspension :— 

2SO a |g Ba + SC 1 2 = BaS s O s + BaCl 2 . 

The free acid, which may be concentrated in a vacuum to a given 
strength, soon decomposes into sulphur, sulphurous anhydride, 
and sulphuric acid in the same manner as the other two acids. 


COMPOUNDS OF SULPHUR AND CHLORINE. 

A chloride of sulphur corresponding to sulphuric anhydride 
(i.e. containing one atom of hexad sulphur united with six atoms 
of chlorine) has not yet been prepared. That consisting of one 
atom of sulphur and four of chlorine—sulphur tetrachloride : 
SC 1 4 —is a very unstable compound. More stable but still easily 
decomposed is sulphur dichoride : SC 1 3 ; while the most stable 
compound of all is 


Bisulphur Bichloride : S 2 C 1 2 . 

This compound is a transparent dark yellow liquid, fuming slightly 
in the air and possessing a powerful piercing odour. Its specific 
gravity is 17, and it boils at 138°. Disulphur dichloride is easily 
obtained by heating sulphur in a retort and passing dry chlorine 
gas over it. The sulphur then melts and the two elements unite 
with one another, the volatile chloride condensing in the cool 
receiver. The raw product is contaminated with sulphur di- 
chloride, from which it may be freed by fractional distillation, 
reserving that portion only which passes over at 138°. Disulphur 
dichloride is a good solvent for sulphur, of which it dissolves more 
than half its weight. It is decomposed by water into sulphur, 
sulphurous anhydride, and hydrochloric acid. 


sulphur Bichloride : SC 1 2 , is a heavy brown oil with similar 
properties to the preceding compound. It is obtained by satuia- 
ting disulphur dichloride with chlorine at the ordinary temperature 
and then removing the excess of chlorine by passing a stream of 
dry carbonic anhydride. The compound begins to boil at 64 , due 
to evolution of free chlorine. It is then decomposed, and partially 



174 


Text-Book of Inorganic Chemistry. 

even at lower temperatures, into disulphur dichloride and chlorine. 
Water slowly decomposes it into sulphur, sulphurous anhydride, 
and hydrochloric acid. Sulphur dichloride forms double compounds 
with some other chlorides— e.g. antimony trichloride, arsenic tri¬ 
chloride, &c. 


Sulphur Tetrachloride : SC 1 4 . 

This is a mobile, pale brown liquid produced by saturating 
either of the other chlorides with chlorine gas at — 22°. The com¬ 
pound is very unstable and breaks up even below o° into sulphur 
dichloride and free chlorine. 

Sulphur also unites chemically with bromine and iodine. The 
most interesting of these compounds, Sulphur hexiodide : SI t ,, 
is deposited by allowing a mixed solution of sulphur and iodine 
in carbon disulphide to gradually evaporate. It is a solid crystal¬ 
line substance, of dark grey colour, with a metallic lustre, is so 
easily decomposed that it gradually loses all its iodine when 
allowed to lie exposed to the air. 


SELENIUM. 

Chemical Symbol : Se .—Atomic Weight: 79. 

This element was discovered by Berzelius in 1817, and on 
account of its similarity to the previously discovered tellurium 
(from tellus, the earth), was named by him selenium (from aeXrjvrj, 
the moon). Berzelius found the new element in the deposit which 
had collected on the floor of a sulphuric acid chamber, fed with 
sulphurous anhydride derived from iron pyrites. Afterwards 
selenium was also discovered in the flue-dust deposited in the 
passages through which the acid vapours had to pass to reach the 
leaden chambers. It was later shown that the element exists, 
combined with various metals—*?.£. lead, silver, and mercury—in the 
mineral kingdom, and just as many samples of iron pyrites contain 
traces of iron selenide as well as iron sulphide, so too some 
varieties of native sulphur contain minute quantities of free 
selenium. Selenium is, however, a rare element, and always 
difficult to obtain in large quantities. 



Selenium. 


175 

Selenium which has been melted and allowed to cool is a dark 
brown amorphous substance with a lustre nearly metallic. It has 
a specific gravity of 4*2, melts at 217 0 , and boils at about 700°. 
Its vapour is of a dark yellow colour, and condenses either as 
bright red flowers of selenium or as a dark metallic-like mass. 
The specific gravity of its vapour "at about 86o° is 7*67, but 
diminishes nearly to that corresponding to the molecule Se 2 
(5*58) at about 1400°. 

If selenium is reduced from a solution of selenious acid ( e.g . by 
sulphurous anhydride) it is deposited as a red amorphous powder. 
This modification is soluble in carbon disulphide, from which it is 
again separated on evaporation as dark-red, transparent crystals, 
isomorphous with monoclinic sulphur. A second crystalline modi¬ 
fication is also known which is insoluble in carbon disulphide, has a 
specific gravity of 4*5, and conducts an electric current feebly; it 
also possesses the remarkable property of conducting better under 
the influence of light than in the dark. This form of selenium is 
prepared by melting the element and keeping the temperature 
constant at 210°, until it has become crystalline in structure. It 
then conducts twice as well in diffused daylight, and nearly ten 
times as well in direct sunlight as in the dark. If the intensity of 
the light is diminished, the conductivity rapidly decreases. 

The extraction of selenium from flue-dust containing it, or from 
its ores, is a tedious process. The various methods chiefly consist 
in oxidizing the selenium to selenious acid and then reducing it 
again by sulphurous anhydride. 

Selenium burns when heated in the air with a bright blue flame- 
to form selenious anhydride. It requires, however, a higher 
temperature for its combustion, and burns more difficultly than 
sulphur. During combustion it gives off an intense and disgusting 
odour, which, however, is not peculiar to selenious anhydride, but is 
probably due to the simultaneous production of small quantities of 
a lower oxide. 


COMPOUNDS OF SELENIUM. 

The compounds of selenium are closely related in their che¬ 
mical properties to those of sulphur. Seleniuretted hydrogen 
corresponds to sulphuretted hydrogen, selenious anhydride to 



176 Text-Book of Inorganic Chemistry. 

sulphurous anhydride, selenic acid to sulphuric acid, and the two 
chlorides of selenium to disulphur dichloride and sulphur tetra¬ 
chloride respectively. But, at the same time, many differences 
may be noted when similar compounds are compared together. 
Sulphurous anhydride, for example, is a gas, and does not unite 
with water to form the corresponding acid, while selenious anhy¬ 
dride is a solid crystalline body and easily forms the corresponding 
selenious acid. Sulphur and sulphurous anhydride are oxidized 
by nitric acid to sulphuric acid, while nitric acid has no action on 
selenious acid and only oxidizes selenium to the lower acid. Sul¬ 
phuric acid can be distilled unchanged, or at least with very slight 
decomposition ; selenic acid breaks up at about 280° into selenious 
acid and oxygen. Sulphur tetrachloride is a very unstable liquid, 
decomposing even below o°; while selenium tetrachloride is a 
crystalline solid which can be heated without decomposition. 

It will thus be seen that considerable differences exist between 
sulphur and selenium with regard to their affinities for other bodies. 
It is clear from the facts mentioned above that selenium has a 
stronger affinity for chlorine and a weaker affinity for oxygen than 
sulphur, and this is confirmed by the fact that sulphur catches fire 
and burns when heated in oxygen or air much more easily than 
selenium. 

Still more remarkable is the affinity of selenium for hydrogen, 
and in this respect again it is at once distinguished from sulphur. 
Sulphur and hydrogen unite directly with one another only imper¬ 
fectly, and the resulting sulphuretted hydrogen is very easily decom¬ 
posed by heat. Selenium, on the other hand, unites with hydrogen 
at about 500°, forming seleniuretted hydrogen, which is only de¬ 
composed into selenium and hydrogen at a high temperature. 

Selenium, with the similar elements tellurium and sulphur and 
the metals of the alkalies, are the only elements which unite directly 
both with hydrogen and oxygen. Thqse elements which possess a 
strong affinity for oxygen— eg. phosphorus, arsenic, silicon, carbon, 
and most of the metals—cannot unite directly with hydrogen ; while 
the few other elements distinguished by their affinity for hydrogen 
—eg. chlorine and palladium—cannot combine directly with oxygen. 



Compounds of Selenium . 


77 




Seleniuretted Hydrogen : H„Se. 

This compound is a colourless gas with an odour resembling 
sulphuretted hydrogen, and is exceedingly poisonous. Minute 
quantities in the air produce painful inflammation of the eyes and 
nasal passages, loss of the sense of smell, dryness in the throat, 
and a species of catarrh, often lasting for weeks together. Great 
care must, therefore, be taken in working with the gas, and in all 
operations during which it may be set free. 

Seleniuretted hydrogen is produced when dry hydrogen gas is 
led over selenium heated to the temperature at which it commences 
to volatilize. It may be obtained in larger quantities by acting 
upon potassium selenide or ferrous selenidewith hydrochloric acich 
Like sulphuretted hydrogen, the gas is soluble in water, but in 
larger quantities. This solution reddens litmus paper, colours the 
skin red from reduced selenium, and becomes oxidized when ex¬ 
posed to the air, selenium being set free. Seleniuretted hydrogen 
precipitates many metals from solutions of their salts as insoluble 
selenides, just as- sulphuretted hydrogen precipitates insoluble 
sulphides. 


Selenious Anhydride : Se0 2 . 

This substance is obtained when selenium is heated in a stream 
of dry oxygen. The selenium burns with a bright blue flame pro¬ 
ducing the anhydride, which is deposited on the cooler portions of 
the tube in long, quadrangular, white needles. It sublimes without 
first melting. Its vapour resembles that of chlorine in colour and 
possesses a piercing acid odour, the disgusting odour produced 
when selenium burns in oxygen or air being probably due to some 
lower oxide. When boiled with water it unites with this substance 
and forms the corresponding — 

Selenious Acid: SeO j 

This dibasic acid is deposited on cooling the hot saturated 
aqueous solution in transparent crystalline prisms resembling those 
of potassium nitrate, and with a strong acid taste. When heated 
these crystals break up into the anhydride and water. The aqueous 
solution when made acid with a few drops of hydrochloric acid 
deposits selenium if warmed with sulphurous anhydride or other 
reducing agents. Particles of dust of organic origin produce the 

N 


178 Text-Book of Inorganic Chemistry. 

same change, so that the lips of glass bottles containing aqueous 
selenious acid become coloured red from reduced selenium. With 
sulphuretted hydrogen selenious acid is decomposed into com¬ 
pounds of selenium and sulphur, water being also formed. 

Selenious acid can also be easily obtained by oxidizing selenium 
with nitric acid ; after the excess of nitric acid has been expelled 
it remains behind as a white crystalline mass. Selenic acid is 
never produced by this reaction. 


c n fOH 

Selenic Acid: Se0 2 iQ^ 

This acid, which corresponds to sulphuric acid in its composi¬ 
tion, and which it resembles in many respects, cannot be obtained 
by oxidizing selenium with nitric acid. It is, however, easily 
produced by leading chlorine into a strong aqueous solution 
of selenious acid, or by the oxidation of the latter substance 
by potassium dichromate. Potassium selenate may also be easily 
prepared by fusing together either selenium or a compound 
containing it with potassium nitrate. An aqueous solution of 
selenic acid may be concentrated until it attains a boiling-point of 
265° and specific gravity of 2-6, but further concentration decom¬ 
poses it into oxygen and selenious acid. The concentrated 
acid is a colourless, oily liquid, which evolves heat when mixed 
with water, and is not affected by sulphurous anhydride or sul¬ 
phuretted hydrogen. It is, however, reduced to selenious acid 
when boiled with hydrochloric acid, chlorine being evolved. 

The selenates closely resemble the sulphates in their properties ; 
barium selenate, like barium sulphate, is insoluble in water and 
dilute acids. 


Selenium Disulphide : SeS 2 . 

This compound separates out as a yellow precipitate when 
sulphuretted hydrogen is led into an aqueous solution of selenious 
acid. The yellow precipitate gradually coagulates to a red mass, 
a colour which it retains on drying. If, on the other hand, 
seleniuretted hydrogen diluted with hydrogen is led into a saturated 
aqueous solution of sulphurous anhydride, the yellow precipitate 
consists chiefly of sulphur diselenide : SSe 2 . Selenium and sul¬ 
phur may be fused together in any proportion. 


Tellurium. 


179 


Chlorides of Selenium, 

Diselenium dichloride : Se^Cl,, is prepared in precisely the 
same manner as disulphur dichloride, which it closely resembles. 

Selenium tetrachloride : SeCl 4 , which is obtained by saturating 
the preceding compound with chlorine or by distilling selenious 
anhydride with phosphorus pentachloride, is a solid body, crystal¬ 
lizing in colourless cubes. When heated it sublimes without 
melting. It dissolves in water and then gradually decomposes 
into selenious and hydrochloric acids. 


TELLURIUM. 

Chemical Symbol : Te .—Atomic Weight : 128. 

This extremely rare element was discovered by Muller v. 
Reichenstein in 1782 in gold ores obtained from Transylvania, but 
was only exactly studied by Klaproth in 1798, and later by Berzelius 
in 1832. 

It occurs native in the mineral kingdom, but more commonly 
combined with gold, silver, lead, and bismuth. The mineral 
tetradymite (bismuth telluride) contains about 50 per cent, of 
tellurium. The ores are found in Hungary and Transylvania, and 
recently in large quantities at several places in North America. 

Tellurium possesses many of the properties of a metal. It has 
a perfect metallic lustre, is nearly silver white in colour, crystalline 
in structure (crystallizing in rhombohedra), is brittle, and therefore 
easily reduced to powder. It conducts heat well, and electricity to 
some extent. Tellurium melts at about 500°, and can be sublimed at 
a higher temperature, producing a bright yellow vapour * its specific 
gravity is 6*4. Carbon disulphide does not dissolve it, but it is 
soluble in warm concentrated sulphuric acid, forming a magnificent 
red solution, from which it is again precipitated unchanged as 
a grey powder when the solution is diluted with water. When 
strongly heated in the air it burns with a brown flame, bordered 
with green, forming tellurous anhydride. Hot nitric acid converts 
it into tellurous acid. 

Like selenium, tellurium unites directly with hydrogen when 
heated in a stream of the dry gas, and then produces telluretted 
hydrogen : H 3 Te. The same substance is also easily obtained by 

n 2 



180 Text-Book of Inorganic Chemistry. 

acting on zinc telluride with hydrochloric acid. It is a colourless 
gas with a disgusting odour, similar in its properties and reactions 
to sulphuretted hydrogen and seleniuretted hydrogen, though not 
so poisonous as the latter. 

The other compounds of tellurium —tellurous anhydride and 

acid : Te 0 2 and TeO |q[], telluric acid : Te 0 2 1 Qpp tellurium 

di- and tetrachloride : TeCl 2 , and TeClj, tellurium disulphide : TeS 2 
—are closely allied to the corresponding compounds of selenium. 
The tellurates are, however, very different in their properties from 
the sulphates and selenates ; those of the heavy metals being 
mostly insoluble in water, while barium tellurate is only difficultly 
soluble in water and dissolves easily in dilute hydrochloric acid. 
Double salts of telluric acid corresponding to the alums have not 
been prepared. 


ELEMENTS OF THE NITROGEN GROUP. 

This group consists of the elements nitrogen, phosphorus, 
arsenic, and antimony. These elements are so different from one 
another both in their physical properties and their chemical beha¬ 
viour, as far as the strength of their affinity for other elements is 
concerned, that it is not at first sight evident why they are placed 
together in a group. Nitrogen does not catch fire when heated in 
the air, and it combines so feebly with other substances that it was 
thought formerly not to unite directly with any other body. Phos¬ 
phorus, on the other hand, is an easily fusible solid, distinguished 
by its strong attraction for oxygen, chlorine, and other elements. 
Arsenic and antimony possess the general physical properties of 
the metals ; they are heavy solids, with a metallic lustre. We shall, 
however, include these two last-named elements under the non- 
metals, because their chemical compounds with oxygen, hydrogen, 
and other elements are closely related to those of nitrogen and 
phosphorus. And this is also the reason why these four so different 
elements are included in one group. In most of their compounds 
the elements of the nitrogen group possess a triad or pentad 
atomicity. 



Nitrogen . 


181 


NITROGEN. 

Chemical Symbol : N .-—Atomic Weight 114. 

Nitrogen occurs free in atmospheric air, of which it is the chief 
constituent; 100 volumes of common air contain about 79 volumes 
of nitrogen and 21 volumes of oxygen. It also occurs in nature 
combined with hydrogen as ammonia, and with hydrogen and 
oxygen as nitric and nitrous acids. From these compounds nitrogen 
enters into the structure of plants and then of animals. It is a 
constituent of albumen, casein, blood, and numerous other organic 
products; Urea, a substance present in large quantities in urine, 
contains a considerable proportion of nitrogen. 

Nitrogen is a colourless gas, without taste or smell, and rather 
lighter than common air. Its specific gravity is exactly 0-972, 
which corresponds to a molecular weight of 0-972 x 28-88 = 28*07, 
the molecule containing therefore two atoms (N 2 ). The gas is only 
slightly soluble in water. Nitrogen, like hydrogen and oxygen, was 
formerly called a permanent gas, as up to quite recently it had not 
been reduced to the liquid form. It has, however, been proved that 
nitrogen can be condensed to a liquid if the cold and pressure to 
which it is exposed are sufficiently intense. 

Nitrogen gas may easily be obtained from common air, by 
abstracting the oxygen which it contains. This may be done by 
burning a piece of phosphorus in a bell-jar full of air and standing 
over water. All the oxygen then unites with the phosphorus, pro¬ 
ducing white vapours of phosphoric anhydride, which afterwards 
dissolve in the water. As the air cools the water rises in the bell- 
jar to supply the place of the consumed oxygen, and when the air 
has become clear it consists almost entirely of nitrogen. A better 
method, especially when a stream of nitrogen is required, is to pass 
atmospheric air, which has been previously dried and freed from 
carbonic acid, over red-hot metallic copper. Finely divided copper, 
such as copper turnings, is heated to redness in a tube of hard 
glass while a slow current of the purified air is passed over it. The 
gas which then escapes at the other end of the tube is pure nitrogen, 
all the oxygen having combined with the copper to form black 
copper oxide. 

Another method for the preparation of nitrogen consists in 


182 Text-Book of Inorganic Chemistry. 

heating ammonium nitrite. This salt then breaks up into nitrogen 
gas and water :— 

NO- 0 (NH 4 ) = N 2 + 2 H 2 0 . 

Finally, nitrogen is also set free when chlorine gas is led into 
strong ammonia. In this reaction, hydrochloric acid is produced, 
which unites with the excess of ammonia, forming ammonium 
chloride :— 

8 NH 3 + 3 C 1 2 = N 2 + 6 NH 4 C 1 . 

If, however, the chlorine is passed until no more free ammonia 
is present, the gas then begins to decompose the ammonium chloride, 
and oily drops of nitrogen chloride, a highly explosive substance, 
are produced. This process cannot therefore be recommended for 
the preparation of nitrogen. 

Free nitrogen possesses so weak affinities for other elements 
that it can only be made to combine directly with a very few of 
them. With oxygen it only unites at the highest temperatures and 
in the presence of water vapour, then producing nitric acid with 
other lower oxides of nitrogen. Flashes of lightning passing 
through moist air cause the oxygen and nitrogen to unite and form 
nitric acid, which can be easily detected in rain water after a 
thunderstorm. 

The direct combination of nitrogen with hydrogen, chlorine, or 
sulphur, has not yet been effected, and only a very few of the metals 
combine directly with nitrogen when heated in the gas. 

Nitrogen neither supports ordinary combustion nor respiration. 
A candle immersed in the gas is at once extinguished. At the 
same time it is not an actual poison, as at once follows from the fact 
that we respire a mixture of nitrogen and oxygen in common air. 
An animal when brought into an atmosphere of nitrogen is simply 
suffocated from lack of oxygen. 


COMPOUNDS OF NITROGEN AND HYDROGEN. 

We have already stated that all attempts to unite nitrogen and 
hydrogen directly with one another have failed. They unite, how¬ 
ever, when the two elements are brought into contact with one 
another in the nascent state. 

The only compound of nitrogen and hydrogen which can exist 



A mmonia. 


183 


in the free state is ammonia. This substance consists of one atom 
of nitrogen united with three atoms of hydrogen, and in which, 
therefore, the nitrogen plays the part of a triad element. 

No compound analogous to hydrochloric acid— i.e. containing 
one atom of nitrogen united with one atom of hydrogen, nor one 
analogous to water, containing two'atoms of hydrogen and one of 
nitrogen—appears to exist in the free state. Both are, however, 
known in many compounds as hypothetical radicals— i.e. unsatu¬ 
rated compounds of triad nitrogen. The former (NH) is called 
imidogen , and the latter (NH 2 ) amidogen. 

Finally, the compound of pentad nitrogen with four atoms of 
hydrogen (NH 4 ), called ammonium , is unknown in the free state, 
but exists in a large number of compounds very similar to those of 
the metal potassium. 


AMMONIA. 

Composition : NH 3 . 

This compound of nitrogen and hydrogen occurs free in nature 
and is produced during the putrefaction of nitrogenous organic 
matter in the absence of strong bases. It is formed, for example, 
in large quantities in putrefying urine, one constituent of which 
(urea) then combining with water and breaking up into carbonic 
anhydride and ammonia. 

Ammonia is a colourless gas with a strong, piercing odour, 
bringing tears to the eyes. It is extremely soluble in water, 1 
volume at the ordinary temperature dissolving more than 800 
volumes of the gas. This solution, the liquor ajnmonice or spirits 
of hartshorn of the shops, possesses the odour of the gas, and 
reacts powerfully alkaline— i.e. turns red litmus paper blue. The 
gas is considerably lighter than air, its exact specific gravity being 
0*589. Its molecular weight is thus 0*589 x 28*88= 17*01, and its 
molecule is represented by the formula NH 3 . 

Under a pressure of seven atmospheres at the ordinary tempera¬ 
ture, or at -40° under the ordinary pressure, dry ammonia gas 
becomes converted into a colourless liquid which freezes to a 
crystalline solid at about - 8o°. 

Ammonia is prepared from the ammonium salts (compounds 
formed by the union of ammonia with acids), and generally from 



184 Text-Book of Inorganic Chemistry . 

ammonium chloride or sal-ammoniac. This salt, which can be 
obtained from ammonia and hydrochloric acid, is decomposed 
again into these constituents when intimately mixed with strong 
bases, such as caustic potash, caustic soda or quick-lime, and 
gently heated. Even when dry ammonium chloride and dry quick¬ 
lime are mixed together, ammonia begins to be given off without 
the mass becoming liquid, whence it follows that the old chemical 
rule, only liquids act upon one another , is not to be taken literally. 
The reaction goes on according to the equation :— 

2 NH 4 C 1 + CaO = CaCl 2 + H 2 0 + 2NH3. 

Calcium chloride and water are thus simultaneously produced, and 
a portion of the latter passes over with the gas on heating. 

In practice, ammonium chloride and quick- or slaked-lime, both 
finely powdered, are intimately mixed together, and the mixture 
brought into a glass flask or an iron vessel if large quantities are 
employed. Heat is then applied and the ammonia dried by passing 
it through a tube containing pieces of quick-lime. It is then either 
collected over mercury or by upward displacement. The sub¬ 
stances generally employed to dry gases—sulphuric acid or calcium 
chloride—cannot be used for ammonia, as both these bodies absorb 
large quantities of the gas. 

If an aqueous solution of the gas is required, it is of course 
unnecessary to dry it. It is then sufficient to separate the solid 
particles which may be carried over with the gas. This is best 
done by passing the gas through a small quantity of concentrated 
aqueous ammonia contained in a wash-bottle. Large quantities of 
heat are evolved when ammonia is absorbed by water, principally 
due to the latent heat set free when the substance is changed from 
the gaseous to the liquid form. Care must therefore be taken 
to cool the vessel containing the water during the absorption of 
the gas. 

The production of ammonia from nitric acid by the action 
of nascent hydrogen is interesting theoretically. This reduction 
may be produced in several ways; for example, by the action of 
dilute nitric acid on zinc. No evolution of hydrogen is noticed, 
as when zinc is dissolved in sulphuric acid, but the whole of the 
hydrogen is used to reduce the nitric acid. The ammonia which 
is formed unites with the excess of acid, producing ammonium 
nitrate ; the zinc is also converted into its nitrate, and water is 
set free:— 


A minonia. 


185 


2NOj-OH + 4H, - N 0 2 - 0 (NH 4 ). + 3H 2 0. 

On adding an excess of caustic soda to the solution, the ammonia is 
again liberated, and may be easily recognized by its odour. 

In order to produce small quantities of liquid ammonia, advan¬ 
tage is taken of the fact that many metallic salts— e.g. calcium 
chloride, silver chloride—absorb ammonia in large quantities and 
liberate the gas again on warming. Dry silver chloride saturated 
with dry ammonia is placed in the closed end, zz, of a thick walled 
bent glass tube, and the open end, c, is then 
melted together (fig. 47). If now the limb c is 
placed in a freezing mixture, and a gently 
warmed, gaseous ammonia is liberated in large 
quantities from the silver chloride and is con- c 
densed to a colourless liquid in the cold part 
of the tube. When the tube cools again, the silver chloride 
reabsorbs the whole of the ammonia, and the experiment may 
therefore be repeated several times with the same tube. 

Although ammonia contains a large quantity of hydrogen, it 
does not catch fire in the air, but in oxygen gas it burns with a 
pale yellowish flame, forming water and free nitrogen. At a high 
temperature it is decomposed into its constituents, hydrogen and 
nitrogen. If the gas is led through a porcelain tube filled with 
broken fragments of the same and heated to bright redness, the 
volume of the gas becomes doubled 

2NH3 = N 2 + 3H 2 . 

4 vols. 2 vols. 6 vols. 

The same decomposition occurs when electric sparks are 
allowed to pass through the gas for a long time. 

Ammonia combines readily with acids and produces real salts, 
which closely resemble the corresponding compounds of potassium. 
These are the ammonium salts and contain the monad radical 
ammonium (NH 4 ), which plays the part of a metal. When, for 
example, ammonia unites with hydrochloric acid :— 

NH 3 + HC 1 = (NH 4 )C 1 , 

we consider that the hydrogen of the acid first unites with the 
nitrogen, and that the ammonium so produced then forms with the 
chlorine the haloid salt. In a similar manner ammonia also unites 
with the oxy-acids, but in these compounds the ammonium is 
united to the acid radical by an atom of oxygen, thus 


b 



Fig. 47 * 


i86 


Text-Book of Inorganic Chemistry. 


no 2 -oh 

+ 

NH 3 = N0 2 -0(NHJ 

Nitric acid 


Ammonium 



nitrate 

so l 0H 
^2 (OH 

+ 

2 NH -sol°( NH <) 
21 NH 3 - SU 1|0(NHJ 

Sulphuric 


Ammonium 

acid 


sulphate 

so l 0H 
JU2 l0H 

+ 

NTH _ QO i OH 

,\n 3 ~ ^j 0 (NH 4 ). 

Sulphuric 


Acid ammonium 

acid 


sulphate 


Ammonia also combines with the acid anhydrides, in which the 
hydrogen necessary to convert it into ammonium is absent. Two 
molecules of ammonia are then decomposed into one molecule of 
ammonium and one of amidogen :— 

2NH3 = NH 4 + NH 2 , 

_ which unite with the anhydride (eg. C 0 2 or SO s ) to form ammonium 
compounds of monobasic amido-acids 

CO + ?NH - CO j NH 2 
c,w 2 + zi\n 3 - ^^|oNH 4 

Ammonium 

carbamate 

In all cases when gaseous ammonia comes into contact with a 
gaseous acid, heavy white fumes are produced of the ammonium 
salt. This affords a ready test for small quantities of free am¬ 
monia. If a glass rod dipped in concentrated hydrochloric acid is 
brought near the surface of a liquid containing free ammonia, white 
fumes of ammonium chloride are at once produced. If the am¬ 
monia is combined with an acid, as an ammonium salt, it must 
first be set free by warming with caustic soda or potash. 

Chlorine at once decomposes ammonia, forming hydrochloric 
acid, which then combines with the excess of ammonia to form 
ammonium chloride (p. 182). 


HYDROXYLAMINE—(OXYAMMONIA). 
Composition : N H 2 • O H. 

In hydrogen peroxide we have already become acquainted with 
a monad radical—hydroxyl—having the composition HO (p. 90). 
We meet with this hydroxyl in a large number of compounds, par¬ 
ticularly in organic chemistry, in which it displaces other monad 



Hydroxylamine—(Oxy ammonia). 187 

elements or radicals, especially hydrogen. And we may look upon 
hydrogen peroxide as water in which one atom of hydrogen is dis¬ 
placed by hydroxyl, thus : O j 

In 

In precisely the same manner, hydroxyl can also displace one 
of the atoms of hydrogen in ammonia, and so form the substance 
hydroxylamine, or oxyammonia. 

Oxyammonia has not yet been prepared in the pure state ; it 
is only known in aqueous solution and in combination with acids. 
It possesses similar properties to those of ammonia, unites with 
acids producing the oxyammonium salts, but is less volatile. Its 
aqueous solution is colourless and reacts strongly alkaline. 

Oxyammonia is easily produced by the reduction of nitric oxide 
or higher oxides of nitrogen with nascent hydrogen. For this 
purpose nitric oxide is led into a mixture of granulated tin and 
hydrochloric acid, to which a few drops of platinic chloride have 
been added to accelerate the evolution of hydrogen. The acid 
liquid then contains stannous chloride together with oxyammonium 
chloride ; it is freed from tin by passing a stream of sulphuretted 
hydrogen through it, and the filtered liquid evaporated to dryness. 
The dry mass is then extracted with absolute alcohol, which dis¬ 
solves the whole of the oxyammonium chloride, together with a small 
quantity of ammonium chloride (easily removed by a few drops of 
platinic chloride). The clear alcoholic liquid on evaporation then 
yields colourless crystals of oxyammonium chloride : NH 3 (OH)Cl. 

From the chloride the sulphate : S 0 2 ( 0 NH 3 0 H) 2 , is easily 
prepared by adding to it the requisite quantity of dilute sulphuric 
acid and evaporating off the hydrochloric acid thus set free. Free 
oxyammonia may then be obtained by adding baryta water to the 
aqueous solution of the sulphate as long as a turbidity of barium 
sulphate is produced. .The filtered alkaline solution then contains 
free oxyammonia. 

If this solution is distilled a portion of the oxyammonia passes 
over with the water vapour, but the greater part is decomposed 
into ammonia, nitrogen, and water :— 

3 NH 2 -OH = NH 3 + N 2 + 3 H 2 0 . 

The aqueous solution of oxyammonia powerfully reduces many 
compounds. Silver and mercury are precipitated from the solutions 
of their salts in the metallic state, and with copper sulphate it gives 
an orange-coloured precipitate of cuprous oxide. 


138 Text-Book of Inorganic Chemistry . 

Oxyammonia is also produced when some metals, especially 
zinc, are acted upon by sulphuric acid in the presence of nitric acid. 
The nascent hydrogen which is then set free reduces a part of the 
nitric acid to ammonia and oxyammonia. If zinc is used the solution 
contains the nitrates of zinc, ammonium, and oxyammonium, and 
on the addition of excess of caustic potash and a few drops of copper 
sulphate gives an orange-coloured precipitate of cuprous oxide. 

The oxyammonium salts, like those of ammonium, are all solu¬ 
ble in water, and can easily be obtained in the crystalline form. 


AMIDOGEN: NH 2 ; IMIDOGEN : NH. 

These are two unsaturated compounds of triad nitrogen, the 
first being a monad, and the second a dyad radical. Both occur 
principally in organic compounds, but amidogen is also known in 
some inorganic compounds ; for example, in the amido-acids. 


AMMONIUM : NH 4 . 

This unsaturated compound of pentad nitrogen has not yet 
been prepared in the free state. It is the radical of the ammonium 
salts, but all attempts to separate it, or rather its molecule : (NH 4 ) 2 , 
have resulted in a mixture of ammonia and hydrogen. 

It is very probable that, should this body ever be obtained in 
the free state, it would be found to possess the properties of a 
metal. N ot only have its compounds the greatest similarity with 
those of potassium and sodium, but it can also unite in the nascent 
state with mercury to form an amalgam—a property peculiar to the 
metals. Ammonium amalgam, with metallic lustre and other re¬ 
markable properties, is easily obtained by adding sodium amalgam 
to a warm, moderately concentrated solution of ammonium chloride. 
The sodium amalgam immediately swells up to a large extent, at 
the same time retaining its metallic lustre, and the mass becomes 
so light that it floats on the surface of the liquid, notwithstanding 
the mercury which it contains. 

The reaction simply consists in the substitution of sodium for 
ammonium in the ammonium chloride, forming sodium chloride, 
and setting the ammonium free. The amalgam is very unstable, 
breaking up immediately into ammonia, hydrogen, and mercury, 




Compounds of Nitrogen and Oxygen. 189 

the first-named substance being at once recognized by its charac¬ 
teristic odour. 

The low specific gravity of the ammonium amalgam makes it 
probable that if the ammonium itself were separated, it would be 
a light liquid or a gas. 


COMPOUNDS OF NITROGEN AND OXYGEN. 

Nitrogen and oxygen unite in five different proportions, and 
produce as many different compounds. These compounds have 
the following names and symbols. Two of them (nitrous anhy¬ 
dride and nitric anhydride), unite with water and form the corie- 
sponding acids :— 


Nitrous oxide . 

. n 2 o 

Nitric oxide 

. NO 

Nitrous anhydride . 

. n 2 o 3 

Nitric peroxide 

. n 2 o 4 

Nitric anhydride 

. N 2 0 5 


Nitrous acid 


Nitric acid 


NO-OH 


N 0 2 -OH 


NITRIC ACID. 

Composition'. HN 0 3 = NO- 2 *OH. 

Nitric acid occurs free in the atmosphere, especially after 
thunderstorms, and is produced directly from its elements by the 
high temperature of the flashes of lightning. Its salts are also 
produced in nature, when nitrogenous organic substances putrefy 
in contact with strong bases (e.g. potash, soda, lime). 

Nitric acid is a colourless, strongly acid liquid, fuming in the 
air. It has a specific gravity of 1-530, boils at 86°, and becomes 
solid at - 50°; it mixes with water, with evolution of heat, in all pro¬ 
portions. The acid is not very stable, and is even partially decom¬ 
posed by light into oxygen and lower oxides of nitrogen, which 
impart a yellow tint to it. It destroys the skin and other animal 
substances and colours them yellow, and, in consequence of the 
ease with which it breaks up into oxygen and the lower oxides of 
nitrogen, is an excellent oxidizing agent, especially for the metals 
Nitric acid, as we have previously remarked (p. 182), is produced 
directly from its elements when electric sparks pass through moist 
air, or when a mixture of hydrogen and excess of oxygen, with 






19 ° Text-Book of Inorganic Chemistry . 

traces of nitrogen, is exploded, or when hydrogen burns in oxygen 
mixed with a small quantity of nitrogen. 

The acid is always obtained from one of its salts—potassium or 
sodium nitrate—by decomposing it with sulphuric acid. On the 
small scale pure nitre (potassium nitrate) is powdered and intro¬ 
duced into a retort furnished with a cooled receiver. It is then 
mixed with concentrated sulphuric acid and gently heated, when 
the whole melts to a homogeneous liquid. As the temperature 
rises, the nearly colourless acid distils over, and is condensed in 
the receiver. It is best to employ one molecule of nitre (ioi), to 
one molecule of sulphuric acid (98)— i.e. about equal parts by 
weight. The reaction is then represented by the following equa¬ 
tion :— 

NOj'OK - SO s {gH = NO,.OH + SO,{°£ 

Potassium Acid potassium 

nitrate sulphate 

If we employ twice as much nitre so as to produce the normal 
sulphate, the first stage of the reaction goes on as before, but the 
acid potassium sulphate which remains behind is then mixed with 
one-half of the nitre employed. At a higher temperature these two 
substances act upon one another, forming another molecule of 
nitric acid and normal potassium sulphate :_ 

S ° 2 (OK + n °2'OK = S 0 2 j°| + N 0 2 - 0 H. 

Acid potassium Normal potassium 

sul P hate sulphate 

The temperature at which this takes place is so high that the 
nitric acid undergoes a partial decomposition into oxygen, nitric 
peroxide, and water. The oxygen passes off, but the nitric peroxide, 
a dark red gas, is mostly absorbed by the nitric acid. The acid 
then becomes of a reddish colour, and gives off brown vapours 
This acid is known as fuming nitric acid. In order to obtain a 
pure acid, it is best to take a small excess of sulphuric acid, and 
then redistil the acid with an equal volume of concentrated 
sulphuric acid. The strong attraction of sulphuric acid for water 
then frees the nitric acid of this impurity. 

Commercial nitric acid is now often prepared from Chili salt¬ 
petre (Nals 0 3 ), this salt being much cheaper than ordinary nitre, 
and yielding weight for weight a greater quantity of nitric acid. It 
is usually coloured yellow, and is often contaminated with chlorine, 


Nitric Acid. 


191 


iodine, sodium sulphate, and iron. These impurities may be readily 
separated by repeated distillation, neglecting the first and last 
portions which come over. It always contains water, the quantity 
of which may be determined by means of an hydrometer. The 
greater the quantity of water the lower the specific gravity. The 
specific gravity of this commerical acid is often as low as 1*33, and 
then contains only about 50 per cent, of the pure acid. 

The specific gravity of the pure acid is, as we have stated 
above, 1*530. An acid which has a specific gravity of 1*424, and 
which boils at about 120°, contains nearly 70 per cent, of pure 
acid. The dilute acid, of specific gravity 1*2, which is employed 
in the laboratory for general purposes, contains only 32 per cent, 
of the pure acid. 

If concentrated nitric acid containing more than 70 per cent, of 
the pure acid is boiled under the ordinary atmospheric pressure, 
the portion which distils over contains less water than that in the 
retort until the boiling-point rises to about 120°, when the compo¬ 
sition of the distillate and of the residue in the retort remains 
nearly constant. This acid of constant boiling-point has a specific 
gravity of 1*42, and contains about 70 per cent, of pure acid. 
Further, if dilute nitric acid containing less acid than 70 per cent, is » 
distilled, the portion which passes over is at first more dilute than 
that in the retort, and rises in strength until it contains about 70 
per cent, of pure acid, when the composition again remains nearly 
constant. It must not, however, be supposed that this particular 
mixture is a chemical compound of water and nitric acid. That 
it is not so is proved by the fact that if the pressure is changed the 
composition of the distillate also changes. 

Nitric acid is a solvent for many substances which are insoluble 
in water—particularly for the metals. Among the few metals 
which withstand its action are gold and platinum. Silver is at 
once attacked by it and converted into the soluble nitrate, a 
portion of the nitric acid being reduced to nitric oxide, producing 
red fumes of nitric peroxide in the air. The reaction goes on 
according to the following equation, which also represents the 
oxidation of many other metals by nitric acid 

4 N 0 2 - 0 H + 3Ag 7 3 N 0 2 - 0 Ag + 2 H 2 0 + NO. 

In consequence of this property of nitric acid to dissolve silver but 
not gold, nitric acid has been long employed to separate the two 
metals from one another. 


192 Text-Book of Inorganic Chemistry. 

Nitric acid is a monobasic acid, as it only contains one atom 
of displaceable hydrogen ; it, therefore, only produces normal 
salts. The nitrates are all soluble in water, and we cannot therefore 
recognize them, as we can the sulphates, by the production of one 
which is insoluble in water. All nitrates are decomposed by strong 
sulphuric acid yielding nitric acid ; and if a solution of a nitrate is 
mixed with sulphuric acid and copper turnings and then warmed, 
the nitric acid which is set free acts on the copper and produces 
nitric oxide, which at once gives rise to brown fumes of nitric 
peroxide and nitrous anhydride on coming into contact with the 
oxygen of the air. Another delicate test for a nitrate is mentioned 
under nitric oxide. 


Nitric Anhydride : N 2 0 5 . 

This substance, as its name implies, is nitric acid minus the ele¬ 
ments of water, and may be considered as two molecules of nitric 

peroxide united by an atom of oxygen, thus ^q 2 j O- 

It is obtained in colourless lustrous crystals belonging to the 
rhombic system when finely-powdered dry silver nitrate is dropped 
into a flask containing dry chlorine. The two substances then 
react upon one another, producing nitric anhydride, silver chloride, 
and oxygen. On opening the flask afterwards the oxygen and 
excess of chlorine are got rid of by blowing in dry air, and the 
anhydride, which melts at 30°, is then easily separated from the 
silver chloride by gentle heating. If heated slightly above its 
boiling point (45 0 ), xt decomposes with explosion into nitric 
peroxide and oxygen, and the decomposition also goes on slowly 
in sealed tubes. It readily dissolves in water, evolving heat, and 
producing nitric acid :— 

^q 2 Jo + H 2 0 = 2NCVOH. 


NITRIC OXIDE. 

Compositio?i : NO. 

This oxide of nitrogen is a colourless gas, slightly soluble in water, 
and with a specific gravity of 1*039. This density corresponds to 
a molecular weight of 1*039 * 28*88 -30, whence its composition. 



Nitric Oxide. 


193 

Nitric oxide possesses none of the acid properties of nitric 
acid. It is distinguished from all other gaseous bodies by its 
remarkable property of combining directly with oxygen at the 
ordinary temperature to form a dark red gas. From this it follows 
that nitric oxide may at once be recognized by the red colour 
produced when it comes into contact with common air. This red 
gas is either nitrous anhydride or nitric peroxide, according to the 
quantity of oxygen present. 

Nitric oxide is obtained from nitric acid by abstracting a portion 
of its oxygen by copper or some other similar metal. Small 
pieces of copper are placed in a flask and then dilute nitric acid, 
of specific gravity 1-23, poured over them. The gas soon begins 
to be given off, and the mixture evolves large quantities of heat, 
and if the flask is not cooled from time to time by dipping it into 
cold water, the reaction may become so violent that the liquid 
froths over. The nitric oxide which is first produced unites at 
once with the oxygen of the air in the flask and forms red fumes 
of nitrous anhydride, but as this compound is absorbed by water, 
the gas when collected in a jar over water soon becomes colour¬ 
less. After a time, when the whole of the air has been expelled 
from the flask, pure colourless nitric oxide is given off The reaction 
may be considered as taking place in two stages. In the first 
the copper is oxidized to copper oxide, and the nitric acid reduced 
to nitric oxide and water ; while, in the second, the copper oxide 
unites with the excess of nitric acid to form copper nitrate. The 
entire reaction is expressed in the following equation :— 

3 Cu + 8NCVOH = 3 no*:o} Cu + zN0 + 4H s O. 

If the mixture of copper and nitric acid becomes too hot the 
nitric oxide may easily become contaminated with nitrous oxide. 
Instead of copper we may also use mercury or silver for the 
preparation of nitric oxide ; but all metals cannot be employed, 
some give off nitrous anhydride or nitric peroxide. 

In nitric oxide the constituents are firmly united together, 
much more firmly than in nitrous oxide, which contains twice as 
much oxygen. A glowing chip of wood, faintly glowing charcoal, 
or burning sulphur is at once extinguished in the gas, but strongly 
glowing charcoal continues to burn brilliantly. Phosphorus can 
be melted in the gas without catching fire, but ignited phos¬ 
phorus goes on burning. In all cases when sub.^tances burn in 

O 


194 Text-Book of Inorganic Chemistry. 

nitric oxide the temperature of combustion must be high enough to 
decompose the gas into its constituents. Among the more easily 
combustible substances, the volatile liquid carbon disulphide is 
distinguished by the readiness with which it burns when mixed with 
nitric oxide. A small quantity of the disulphide is poured into a jar 
of the gas closed with a glass plate, the jar slightly agitated so as 
to volatilize the disulphide, and the mixture then ignited. The 
carbon disulphide then burns with so brilliant a flame that it can 
scarcely be borne by the naked eye. 

Strong as well as dilute nitric acid dissolves considerable 
quantities of nitric oxide and becomes coloured brown, yellow, 
green, or blue according to the concentration of the acid. A 
solution in the concentrated acid (specific gravity 1*35) is brown, 
in an acid of specific gravity 1*25 is blue. On heating the acid 
the nitric oxide is again given off. 

Nitric oxide is also absorbed by concentrated solutions of 
ferrous sulphate, yielding a solution of a dark brown colour. This 
property of ferrous sulphate is used to detect small quantities of 
nitric acid (free or combined) contained in a solution. To the 
solution which is to be tested for nitric acid or a nitrate a little 
concentrated sulphuric acid is added, and then a crystal of ferrous 
sulphate. The sulphuric acid serves to set the weaker nitric acid 
free, and this is reduced to nitric oxide by the ferrous sulphate ; the 
gas then unites with the excess of ferrous sulphate to form a dark- 
coloured mantle around the crystal. 


NITROUS ANHYDRIDE AND NITRIC PEROXIDE. 

These two oxides of nitrogen, which are intermediate between 
nitric acid and nitric oxide in composition, are so similar in the 
gaseous state, that one can scarcely be distinguished from the other, 
especially as they are generally produced together. Both are 
brown gases, with a strong unpleasant odour, and both are produced 
by the union of nitric oxide with oxygen—when the nitric oxide is 
in excess, nitrous anhydride is formed, and with excess of oxygen 
nitric peroxide. Both gases can be easily condensed to the liquid 
state by lowering the temperature ; liquid nitrous anhydride is 
blue, while liquid nitric peroxide is brown. 

A mixture of nitrous anhydride and nitric peroxide is easily 



Nitrous Anhydride and Nitric Peroxide . 195 

obtained by heating arsenious anhydride, white arsenic (in lumps) 
with nitric acid of specific gravity 1-33. When the reaction has 
once been started it goes on without further heating until finished 
oxygen bemg abstracted from, the nitric acid while the arsenious 
anhydride is oxidized to arsenic .acid. The red gases are led 
t rough a tube surrounded by a freezing mixture and are there 
condensed to a deep blue liquid. This liquid begins to boil at 
+ 2 , and the boiling point gradually rises to 22° It is a mixture 
of the volatile nitrous anhydride, and the less volatile nitric 
peroxide. From this mixture the two substances may be easily 
obtained in the pure state. 


Nitrous Anhydride : N o 0 3 . 

This substance is obtained from the blue liquid mentioned 
above by passing its vapour mixed with nitric oxide through a red- 
hot glass tube. The nitric peroxide which it contains then 
unites directly with the nitric oxide to form nitrous anhydride • 
N 0 2 + NO = N 8 O s , which can be condensed in a freezing mixture 
to an mdigo coloured liquid. This liquid does not solidify at - -o° 
and begins to boil below o°, while at higher temperatures it easily 
decomposes into nitric peroxide and nitric oxide. Nitrous anhy¬ 
dride dissolves in a small quantity of ice-cold water and forms 
Nitrous acid : NO-OH, according to the equation : N o 0 3 + H,Q 
== 2NO-OH. It easily decomposes when slightly warmed into 
nitric acid, nitric oxide, and water :— 

3NO • OH = NO s OH + 2NO + H 0 0 . 

More stable than the acid itself are its compounds—the nitrites 
Potassium nitrite (NO-OK), which contains one atom of oxygen 
less than potassium nitrate (N 0 8 - 0 K), is easily produced by 
fusing the latter substance until the oxygen which is given off 
begins to be mixed with nitrogen. The potassium nitrite^can then 
be extracted from the powdered mass by means of alcohol, in which 
it is soluble. 

If potassium nitrite, or any other salt of nitrous acid, is treated 
with dilute sulphuric acid, it is at once decomposed; the nitrous 
acid which is thus set free imparts a transient blue colour to the 
liquid, and soon breaks up into water and red vapours of the lower 
oxides of nitrogen. An aqueous solution of potassium nitrite 
when mixed with silver nitrate produces a white precipitate of 
silver nitrite, thus :— 


196 Text-Book of Inorganic Chemistry . 

NO-OK + N 0 2 • O Ag = NO-OAg + N 0 2 - 0 K. 

Ammonium nitrite (NO-ONH 4 ), which, it will be remembered, 
breaks up into water and nitrogen when heated (p. 182), is also 
produced in minute quantities when water rapidly evaporates, a 
portion of the water then uniting with the nitrogen of the atmo- 

Sphere: ~ 2 H 2 0 + N 2 - NO.ONH, 

The same compound is also produced when ammonia is oxi¬ 
dized by a mixture of ozone and oxygen 

2NH3 + 0 3 = NO-ONH, + H 2 0 , 
and the salt is then further oxidized to ammonium nitrate. 


Nitric Peroxide : (N 0 2 ) 2 - N 2 0 4 . 

This compound of nitrogen and oxygen possesses all the 
properties of a compound radical, and resembles chlorine in many 
respects. A more suitable name for it than nitric peroxide would 
be nitrvl, corresponding to hydroxyl, sulphuryl, &c. 

Nitric peroxide is best obtained from the crude blue liquid re¬ 
ferred to on the preceding page. When this liquid is kept in a 
freezing mixture and a current of oxygen led through it the nitrous 
anhydride which it contains becomes oxidized, and pure nitric 
peroxide is obtained. Nitric peroxide may also be prepared by 
leading a mixture of nitric oxide and excess of oxygen into a cooled 
tube, or, finally, by carefully heating dried lead nitrate m a retort, 
when the salt suffers the following decomposition 


°1 
NO*• O f 


N0 2‘^ 1 -Pb 


(N 0 2 ) 2 + O + PbO. 


Nitric peroxide is a pale brown liquid, of specific gravity 1-45 ; 
it boils at 22 0 and then forms a brown gas, which becomes of 
a darker colour when heated. It freezes at -2o°to a colourless 
crystalline mass, which again melts at about -12 0 . The vapour 
density of nitric peroxide at low temperatures has been found to be 
2'65 at high temperatures 1*58. The former number approaches 
the density required for the formula N 2 0 4 , which would be 

(2 X 14) + (4 x 16) _ 92 _ = while that of the latter is about 

-“^88 28-88 

one-half of this, or corresponds to the molecule N 0 2 . 

From this we conclude that at low temperatures the molecule of 
nitric peroxide is composed of twice N0 2 or N0 2 - N0 2 , which, when 




197 


Nitrous A nhydride and Nitric Peroxide. 


heated to ioo°—150°, breaks up into separate molecules of the com¬ 
position N 0 2 . The same change probably takes place when many 
elements—^. oxygen or chlorine—are raised to a very high tempe¬ 
rature. The atoms, which at ordinary temperatures are united to 
form molecules, then divide into separate atoms and so acquire a 
far stronger power of affinity than “in the ordinary molecular state. 
In the same way nitric peroxide unites much more readily with other 
bodies at ioo°than at the ordinary temperature ; for example, it can 
then combine directly with chlorine, bromine, cyanogen, &c. 

Concentrated, colourless nitric acid dissolves large quantities of 
nitric peroxide, producing the recj. or fuming nitric acid. This acid 
is a much more powerful oxidizing agent than ordinary nitric acid, 
because the nitric peroxide which it contains parts with a portion of 
its oxygen far more readily than nitric acid itself. 

In contact with cold water nitric peroxide is decomposed into 
nitric and nitrous acids :— 

(N 0 2 ) 2 + H 2 0 = N 0 2 -OH + NO-OH; 

the nitrous acid gradually decomposing into nitric acid, nitric 
oxide, and water (p. 195). 

Chlorine, which at low temperatures is without action on nitric 
peroxide, combines with it when a mixture of the two gases is led 
through a heated tube, producing nitryl chloride : N 0 2 C 1 , a yellow 
liquid boiling at + 5 0 , which is decomposed by water into nitric 
and hydrochloric acids. 

The same compound is also produced by leading a stream of 
chlorine over gently heated silver nitrate 

N 0 2 -OAg + Cl 2 = N 0 2 C 1 + AgCl + O. 


If liquid sulphurous acid and nitric peroxide are heated together 
in sealed tubes, or if nitric oxide is allowed to act upon sulphuric 
anhydride, a crystalline compound is produced, of the composition : 

S 2 0 5 -2N0 2 , or O ||o 2 -N 0 2 ’ andwhich ma Y be considered either 
as the dinitryl of disulphuric acid, or as the anhydride of the next 
following substance, nitrosulphonic acid. This compound melts 
at 217 0 , and is decomposed by water into sulphuric acid, nitric 
oxide, and nitric peroxide. 

Concentrated sulphuric acid absorbs nitric peroxide and pro¬ 


duces white crystalline nitrosulphonic acid\ S0 2 |qjj 2 (p- 166), ii 


198 


Text-Book of Inorganic Chemistry. 

which one atom of hydroxyl is displaced by one of nitric peroxide, 
nitric acid being probably formed at the same time. This com¬ 
pound is also decomposed by water into sulphuric acid, nitric oxide, 
and nitric peroxide. 

Aqua Regia .—This liquid, which owes its name to its property 
of dissolving gold, the king of metals, as well as platinum, is a 
mixture of three parts of hydrochloric acid and one part of nitric 
acid. The two acids act upon one another at the ordinary tem¬ 
perature, the mixture becoming yellow and giving off minute 
bubbles of gas smelling like chlorine. This gas, which is evolved 
in quantity on gently warming, is a mixture of chlorine and 
niirosyl chloride'. NOC 1 . In the nascent state the latter sub¬ 
stance gives up its chlorine to the gold or platinum, produc¬ 
ing.soluble chlorides of these metals. The reaction by which it is 
produced' 4 s shown by the following equation :— 

NCVOH + 3HCI = NOC 1 + 2H a O + Cl 2 . 


NITROUS OXIDE. 

Composition : N 2 0 . 

This substance is a colourless, odourless gas, with a somewhat 
sweet taste, and a density of 1*52, which corresponds to amolecular 
weight of 43*9. Under strong pressure (thirty atmospheres at o°) 
it is condensed to a colourless liquid, which is lighter than water, 
and with which water does not mix. The liquid boils at about 
--92 0 , and when allowed to evaporate, rapidly solidifies to a white 
crystalline solid. Nitrous oxide is only slightly soluble in water, 
and can, therefore, be collected over this liquid. 

The gas is easily prepared by heating crystallized ammonium 
nitrate in a retort. The salt then melts and completely decom¬ 
poses, with much frothing and apparent boiling, into nitrous oxide 
and water. Thus :— 

no 2 onh 4 = n 2 o + 2 h 2 o. 

It is not advisable to decompose the whole of the salt employed, as 
overheating may cause the nitrous oxide to decompose into its 
constituents with an explosion. This method quickly yields large 
quantities of pure nitrous oxide, if care is taken that the ammonium 



Nitrous Oxide. 


199 


nitrate contains no ammonium chloride, otherwise the gas becomes 
contaminated with chlorine compounds. Nitrous oxide is also 
produced when zinc or one of a few other metals is dissolved in very 
dilute nitric acid. The gas obtained in this way is, however, never 
pure, but always contains more or less nitric oxide. 

The most remarkable property of nitrous oxide is the readiness 
with which it allows combustible bodies to burn in it. Combustion 
takes place in the gas almost as energetically as in pure oxygen, 
much more so than in nitric oxide, although this gas contains 
twice as much oxygen. From this we conclude that the nitrogen 
and oxygen are much more loosely combined in nitrous oxide than 
in nitric oxide. In fact, nitrous oxide is decomposed into its con¬ 
stituents even when passed through a red-hot tube. A glowing chip 
of wood catches fire in nitrous oxide almost as easily as in oxygen. 
Phosphorus and charcoal burn brilliantly, and sulphur, if strongly 
* ignited, continues to burn with a bright flame. Feebly burning 
sulphur is extinguished in the gas, because the temperature is not 
high enough to decompose it into its constituents. That nitrous 
oxide, although a chemical compound of nitrogen and oxygen, sup¬ 
ports combustion much better than the mechanical mixture of the 
same two substances which make up common air is partly because 
the compound contains thirty-three per cent, by volume of oxygen, 
while common air has scarcely more than twenty per cent, by 
volume, and is partly due to the fact that when nitrous oxide is de¬ 
composed a considerable quantity of heat is set free. 

Nitrous oxide is chemically an indifferent body, and does not 
enter into combination with any substance. Its peculiar pro¬ 
perty of producing intoxication when inhaled is of physiological 
interest. When breathed in small quantities it produces a pleasur¬ 
able nervous excitation, whence its common name of laughing-gas. 
But if mixed with atmospheric air and inhaled for some time it 
produces, like chloroform, temporary insensibility, and without 
the unpleasant after-effects of the latter substance. It is largely 
used for minor surgical operations, particularly in dentistry, to 
produce insensibility to pain and to allow the surgeon to operate 
more easily. 1 

1 The composition of nitrous oxide and of nitric oxide may be easily 
determined by burning a smali piece of potassium in a given volume of the 
gas, standing over mercury. The gas is thus decomposed, its oxygen uniting 
with the potassium and the nitrogen being set free. Nitrous oxide is then 
found to contain its own volume of oxygen and nitric oxide one-half its volume, 
whence the formulae N 2 0 and NO may be easily deduced.— Ed. 


200 


Text-Book of Inorganic Chemistry. 


ATMOSPHERIC AIR. 


Although the atmosphere is not a chemical compound, it is 
scarcely of less interest and importance to the chemist than water. 
The importance of the atmosphere for chemistry and other allied 
sciences is at once evident when we remember that the air has 
always played a large part in building up and in disintegrating the 
solid crust of the earth, that all processes of combustion or oxida¬ 
tion on the earth’s surface are supported by it, that its constituents 
nourish both plants and animals, and that the state of health of 
individuals as well as of entire communities depends upon the 
purity of the air as well as of the water supplied to them. 

We have already learnt that atmospheric air deprived of its 
other admixtures consists of about 21 volumes of oxygen mixed 
with about 79 volumes of nitrogen. Besides these two gases 
atmospheric air always contains very varying quantities of water 
vapour, small quantities of carbonic acid (3 to 4 volumes per 10,000), 
and traces of ammonia (a few parts per million). The composition 
of atmospheric air by volume, taking the mean quantity of water 
vapour, is then :— 


Nitrogen . 
Oxygen . < 

Water (mean) . 
Carbonic acid . 
Ammonia 


= 78*35 volumes 
= 20*77 „ 

= 0*84 „ 

= 0*04 „ 

= traces. 


ioo*oo volumes. 

These numbers lead to several interesting questions, the first 
of which is whether the nitrogen and oxygen contained in atmo¬ 
spheric air are chemically united or only mechanically mixed. If 
we abstract the water vapour and carbonic acid from the air, the 
proportion of nitrogen and oxygen then becomes :— 

Nitrogen.= 79*04 volumes 

Oxygen.= 20*96 „ 

ioo*oo volumes, 

and it has been found that ever since the composition of the air 
has been accurately known, this proportion has not appreciably 
varied, and that whether the air is taken from over the sea, from 
the surface of the earth inland, or from high mountain tops, even 






201 


A tmospheric A ir. 

as high as 14,000 feet above the sea-level, the proportion of nitrogen 
and oxygen is always almost exactly the same. This speaks 
strongly for the supposition that the two gases are chemically united 
in atmospheric air. 

But in spite of this we know for certain that the gases are only 
mixed together—that they do not form a chemical compound in the 
atmosphere. This conclusion is supported by the following facts :— 

i. If we mix together 21 volumes of oxygen with 79 volumes of 
nitrogen, we obtain a gas which possesses all the properties of 
common air deprived of its water and carbonic acid. But we 
notice no elevation of temperature, which would certainly take place 
had the gases combined together chemically. 

ii. If cold boiled water is shaken up with atmospheric air, the 
water takes up more oxygen than nitrogen, and the air which re¬ 
mains behind is then poorer in oxygen than common air. Or, if 
common air is passed through a porous tube more of the lighter 
nitrogen diffuses through in a given time than the heavier oxygen. 
Neither of these partial separations of the nitrogen and oxygen in 
common air could take place if the two substances were chemically 
united. 

iii. Finally, every chemical compound contains its constituents 
not only in a fixed proportion, but also in weights which are some 
simple multiple of the atomic weight of its constituents. Atmo¬ 
spheric air contains its nitrogen and oxygen in a constant propor¬ 
tion by volume and weight, but not iji simple atomic proportions. 
The composition of air by weight is :— 

Nitrogen. 77 parts 

Oxygen . . • • • 2 3 » 

100 parts, 

and if we divide these numbers by the respective atomic weights 
of the elements, we obtain the ratio :— 

77 : 2 3 = 5 - 55 : r44 = 3-85 : 1, 

14 16 

a ratio which, although it approaches 4 : 1, is far from being a 
simple one, as it would be in a chemical compound. The difference 
in the composition of the air, and of a compound containing four 
atoms of nitrogen and one of oxygen, becomes more apparent if we 
calculate the percentage composition of this unknown body [N 4 Oj. 
We find it to be:— 



202 


Text-Book of Inorganic Chemistry. 


4 N = 56 . 
0 = 16 . 


= 77-8 per cent. 


= 22 ‘2 


72 


IOO'O 


According to which the air contains o-8 per cent, nitrogen less, 
and o-8 per cent, oxygen more than this unknown compound, 
a difference which is much too great to be ascribed to experi¬ 
mental errors, especially as the methods employed for the analysis 
of such simple gaseous mixtures leave nothing to be desired in exact¬ 
ness, and indeed exceed in this respect almost all other analytical 
methods. 

A further and much disputed point is whether the atmospheric 
air has always possessed the same composition which it does at 
present, and whether its composition will not change in the course 
of ages. The occurrence of nitrogen in the mineral kingdom as 
nitre, with that which is contained in the substance of plants and 
animals, is so extremely minute compared with the immense 
quantities of nitrogen in the air that little change has probably 
taken place in the absolute quantity of nitrogen present in the air. 
But the same cannot be said of the oxygen. 

Although we have no means of even approximately estimating 
the temperature of the earth when it was in the liquid state, still it 
is more than probable that many of the oxygen compounds which 
now form the crust of the eart^i were built up from their constituents 
as the earth gradually cooled to the still extremely high tempera¬ 
tures at which the compounds of oxygen with silicon, calcium, 
hydrogen, &c., can exist. Experiment has shown that water is 
not only decomposed by the electric current, but also when 
raised to a very high temperature. If at any time the tempera¬ 
ture of the earth was higher than that requisite to decompose 
water, its atmosphere must have contained, besides nitrogen and 
hydrogen, much larger quantities of oxygen than after it had 
cooled down. Only when the temperature fell low enough could 
the hydrogen and oxygen unite together with enormous explosions 
to form water vapour. Possibly the protuberances of the sun’s 
atmosphere, which are considered to be chiefly glowing hydroeen, 
are produced by the union of immense quantities of hydrogen 
and oxygen at heights in the sun’s atmosphere where the tempera¬ 
ture is low enough to allow these two substances to combine 
with one another. 


203 


A tmospheric A ir. 

Whether the amount of carbonic acid present in the atmosphere 
was ever greater than is now the case, cannot be satisfactorily 
answered in the present state of our knowledge. We know that 
during geological time immense quantities of carbonic acid, as 
calcium carbonate (limestone), have become a portion of the solid 
crust of the earth. A further quantity of carbonic acid has also 
been fixed in the form of our deposits of coal, the remains of extinct 
plants. At the same time it must not be forgotten that the air is 
continually receiving carbonic acid from the interior of the earth 
by volcanoes, active and extinct, and how far these two processes 
balance one another we are at present unable to say. The luxu¬ 
riant vegetation which produced our coal-fields seems to indicate 
that the amount of carbonic acid then present in the atmosphere 
was greater than now. 

The quantity of oxygen contained in the atmosphere is propor¬ 
tionally small, and must be continually diminished by the respira¬ 
tion of men and animals and by the enormous quantities of coal 
which are continually being burnt. At the same time the quantity 
of carbonic acid in the air is increased by these processes. The 
question thus suggests itself, Can this enormous consumption of 
oxygen ever deprive the air of such quantities of this gas as to render 
respiration impossible ? We know the quantity of oxygen which is 
daily consumed by a man, from which we can easily obtain the 
•quantity of oxygen annually consumed by the entire human race 
(reckoned at 1,500,000,000); and we can further calculate the total 
quantity of oxygen contained in the air. Now, if we consider that 
the oxygen consumed by animals, processes of putrefaction and 
combustion, is even nine times more than that consumed by man, 
then in 1,800 years the total quantity of oxygen in the air would 
only be diminished by about one-fifth per cent. But the idea that 
the air can ever become too poor in oxygen for respiration may be 
discarded when we remember the action of green plants on atmo¬ 
spheric carbonic acid. Green plants require carbonic acid for their 
existence and growth as much as animals require oxygen ; under 
the influence of sunlight their green parts decompose the carbonic 
acid, assimilating the carbon and setting the oxygen free. Thus, 
by the mutual action of plants and animals, the proportions of 
oxygen and carbonic acid in the air remain practically unchanged 
for a long time. 

Water and air, once considered to be chemical elements, are 
still important elements in another sense, and especially for the 



204 Text-Book of Inorganic Chemistry. 

science of health. The remains of Roman aqueducts illustrate 
the importance which was attached to a good supply of pure water 
2,000 years ago by a nation without our present scientific know¬ 
ledge. But in this respect we are still far behind the Romans. 
For centuries we have contented ourselves with impure well-water, 
and only in recent years have cholera, typhus, and other epidemics, 
together with a general high death-rate, been attributed to an 
impure and insufficient water supply. 

The question of ventilation, or the supply of pure fresh air, is 
even more neglected. It is well known how detrimental it is for a 
number of persons to remain long in a closed space where the air 
becomes charged with carbonic acid, water vapour, and other 
noxious emanations from the body, and it is also well known how 
easily this great evil may be avoided by leading away the bad air 
and supplying good air in its place. But, notwithstanding this 
knowledge, public and private buildings are continually being 
erected with insufficient ventilation, or, in some cases, with even 
none at all. 


COMPOUNDS OF NITROGEN WITH THE 
HALOGENS. 

Nitrogen combines with the halogens even less easily than with 
hydrogen, and never unites directly with them. The halogen com¬ 
pounds of nitrogen can only be obtained, like ammonia, by indirect 
methods; unlike ammonia, which they resemble in composition, 
they do not unite with acids, and they are further distinguished by 
their great instability. 

The elements in these compounds are so loosely combined, 
that mere contact with some indifferent liquid or solid often suffices 
to decompose them with a powerful explosion. They are therefore 
dangerous substances to deal with, and their properties ought to 
be known before commencing experiments with them. The best 
known of these compounds are nitrogen chloride and nitrogen 
iodide. 

Nitrogen Chloride : NC 1 S (?). This compound, which may be 
considered as ammonia in which the hydrogen is partly or entirely 
displaced by chlorine, but of which the exact composition is un- 



Compounds of Nitrogen with the Halogens. 205 

known, is obtained by the action of free chlorine on ammonia, with 
simultaneous production of hydrochloric acid. 

Nitrogen chloride is a yellow oily liquid, with a specific gravity 
of i*6, and sinks in water. It possesses a powerful, piercing odour, 
by which minute quantities of it may be detected. 

The compound is easily obtained by passing chlorine into a 
saturated solution of ammonium chloride at about 30°, when the 
following reaction probably takes place :— 

NH 4 C 1 + 3CU = 4HCI + NC 1 3 . 

Aqueous ammonia is even more easily decomposed than the 
chloride, but cannot be employed, as the nitrogen chloride is again 
decomposed by free ammonia. 

Owing to the great danger in experimenting with nitrogen chlo¬ 
ride, it is not advisable to prepare large quantities of it. By the 
following method, small quantities, sufficient to observe its odour 
and explosive properties, may be safely obtained. 

In a strong leaden dish about eight inches wide and four inches 
deep is placed a second smaller dish, also of lead, and about three 
inches wide, and the large dish then filled to about two-thirds with 
a luke-warm saturated solution of ammonium chloride. A cylin¬ 
drical glass jar is also filled with the same liquid, inverted in the 
dish and suspended by a wire just over the smaller dish. Chlorine 
gas is now led into the cylinder until the liquid is completely dis¬ 
placed, when the apparatus is allowed to stand quietly. As the 
liquid now gradually rises in the jar, small oily drops of nitrogen 
chloride appear on the surface, which gradually fall inco the small 
leaden dish. When nearly all the chlorine is absorbed, the small 
dish is removed, a portion of the solution poured off, and the com¬ 
pound exploded. The explosion is best brought about by tying 
an oiled feather to the end of a long stick, and then bringing it 
into contact with the nitrogen chloride ; the whole then instantly 
explodes, while the liquid is shot up several feet high and the base 
of the leaden dish pressed quite flat. 

The force of the explosion is so great that a drop of the chlo¬ 
ride of the size of a pea, exploded under water in a porcelain dish 
with an oiled feather, not only breaks the dish into small fragments, 
but forces the piece immediately under the drop into the wooden 

The discoverer of nitrogen chloride, the French chemist Dulong, 
who, being unacquainted with its explosive character, attempted to 


20 6 Text-Book of Inorganic Chemistry. 

distil some in a retort, lost several fingers of one hand by the ex¬ 
plosion. 

If in any way larger quantities of nitrogen chloride have been 
obtained than it is considered advisable to explode, the substance 
may be left to itself under the liquid. After some little time, it will 
gradually disappear owing to spontaneous decomposition. 

Nitrogen chloride is also obtained by the electrolytic decom¬ 
position of ammonium chloride. If two platinum plates are im¬ 
mersed in a warm saturated solution of the salt, and a strong 
current passed for a short time, the chlorine, which is evolved at 
the positive plate, becomes converted into nitrogen chloride. A 
powerful explosion then ensues if the two plates are brought into 
contact. 

Tffitrogren iodide : NI 3 (?). This compound is a dark-brown 
solid, closely resembling precipitated iodine. It is obtained when 
a saturated alcoholic solution of iodine is mixed with excess of 
strong ammonia. The precipitate is then brought on a filter and 
well washed with water. As long as the substance is moist, it 
bears pressure and contact with all those substances which so 
easily explode nitrogen chloride ; but if dried, mere contact with a 
particle of dust or the percussion produced by closing a door is often 
sufficient to decompose it into its constituents with an explosion. 


PHOSPHORUS. 

Chemical Symbol : P .—Atomic Weight'. 31. 

Phosphorus is not found in the free state in nature, but occurs 
in the mineral kingdom principally as calcium phosphate in the 
minerals phosphorite , apatite , &c., and in coprolites. All soils 
contain phosphorus, but usually in only minute quantities. From 
the soil the phosphorus is taken up by plants, which produce very 
complex organic compounds containing it, and is collected by them 
in various parts of their structures, particularly in their seeds. 
These phosphorus compounds then pass into the bodies of animals 
with the nutriment contained in the plants ; they form an impor¬ 
tant constituent of the blood, nerve substance, &c., and undergo an 
oxidation in the body, being again converted into phosphoric acid 



Phosphorus . 207 

In combination with lime, this phosphoric acid forms the greater 
part of the mineral substance of the bones, while a portion is 
expelled from the body in the faeces, and a larger portion in the 
urine, and is thus returned again to the soil. 

Phosphorus was first obtained by its discoverer, Brand, of 
Hamburg (1669), from concentrated urine. A hundred years later 
Scheele recognized the fact that the urine as well as the bones 
contain calcium phosphate, and that it is this substance which 
yields the phosphorus obtained from urine. The name phosphorus 
| is derived from its property of shining in the dark (<££*, light, and 
cpepeiv, to bear). 

We are acquainted with phosphorus in at least two different 
modifications, which are so different in appearance and in their 
physical as well as in some of their chemical properties that they 
might have been thought to be different elements if it were not 
easy to convert one into the other, and if the products of their 
union with other elements were not the same. 

The variety longest known, and commonly called ordinary or 
yellow phosphorus, is, when freshly prepared, a colourless trans¬ 
parent solid, which is coloured yellow under the action of light, and 
has a specific gravity of 1-83. It maybe melted under water at 
a temperature of 44 0 to a colourless liquid, easily solidifying again 
and sometimes assuming a crystalline form. At the ordinary 
temperature phosphorus is as soft as wax, but becomes brittle at 
low temperatures. Phosphorus boils at 290°, and when out of 
contact with air maybe distilled unchanged. Its gas is colourless, 

[ and has a specific gravity (at i,ooo°) of 4-5. From the density of 
its vapour its molecular weight is 4*5 x 28*88 = 129*9, or, in other 
words, the molecule of phosphorus in the gaseous state has the 
j composition P 4 and contains four atoms. 

Phosphorus heated in the air slightly above its melting point 
catches fire and burns, producing white fumes of phosphoric* 
anhydride. At lower temperatures, when exposed to the air it 
oxidizes without the production of flame, forming phosphorous anhy¬ 
dride. The white fumes produced when a stick of phosphorus is 
exposed to the air, and which are luminous in the dark, are due to 
; phosphorous anhydride, and the garlic-like odour noticed at the 
| same time is produced by the same substance. Ordinary phos¬ 
phorus is so inflammable that it catches fire when rubbed on a 
I rough surface ; it should never be held, even for a short time, be¬ 
tween the fingers, as burning phosphorus produces very painful, 




2 oS Text-Book of Inorganic Chemistry. 

slowly healing wounds. This modification of phosphorus is always 
preserved under water, in which it is insoluble. And although it 
possesses a very strong affinity for oxygen, it does not decompose 
water like potassium and sodium. It is slightly soluble in alcohol 
and ether, but easily soluble in carbon disulphide, from whith solu¬ 
tion it again separates in crystals belonging to the regular system. 
Phosphorus in this modification is extremely poisonous ; even small 
quantities have been known to produce fatal effects. 

A second modification of phosphorus is the red or amorphous 
variety. This is a red-brown amorphous powder, of specific 
gravity 2-2 (or heavier than common phosphorus). It does not 
melt, even when heated up to 250°, neither does it produce white 
fumes when exposed to the air, nor is it luminous in the dark. It 
catches fire in the air only when heated up to 260°, is insoluble in 
carbon disulphide, is odourless, and not poisonous. 

This red modification is produced when ordinary phosphorus 
is heated in an atmosphere of carbonic acid or hydrogen to a 
temperature of 240° for some time; the change goes on more 
quickly if the phosphorus is heated in sealed tubes to 300°. But 
in all cases the weight of the phosphorus neither increases nor 
diminishes. If iodine is mixed with an excess of ordinary phos¬ 
phorus and the mixture heated, the portion of phosphorus which does 
not combine with the iodine is converted into the red variety. A 
small quantity of iodine suffices to produce the change in a large 
quantity of phosphorus. The same change is also slowly produced 
by the action of light. Phosphorus exposed to the light becomes first 
yellow and then red. In order to purify the red modification from 
common phosphorus mechanically mixed with it, it may be re¬ 
peatedly shaken with carbon disulphide or boiled with caustic soda. 

It is remarkable that just as ordinary phosphorus is converted 
into the red variety on heating, so also the latter may be changed 
-into the former by the action of heat. Amorphous phosphorus 
heated in a retort filled with carbonic acid to 260°— i.e. to a 
temperature only slightly higher than is required for its production 
—is completely changed into the common form. And in this case, 
too, no change in weight is observed. 

Since phosphorus has been employed for important technical 
purposes, particularly for the manufacture of lucifer matches, large 
quantities are annually manufactured. The material employed 
for its extraction is not urine, from which it was first prepared, but 
is the calcium phosphate of burnt bones, or as one of the several 


209 


Phosphorus. 


forms in which the compound occurs in the mineral kingdom. 
Although the attraction of phosphorus for oxygen is very great, it is 
exceeded by that of glowing charcoal. Phosphoric acid, in the 
proper form, when mixed with powdered charcoal and heated to 
bright redness, is completely reduced with formation of carbonic 
oxide. But if the compound of phosphoric acid and lime, as it 
occurs in nature or in burnt bones (calcium phosphate, Ca 3 (P 0 4 ) 2 ), 
were employed, no reduction would take place. It is only the free 
acid, or compounds of calcium containing more phosphoric acid 
than the normal salt, which are decomposed on heating with 
charcoal. In order to obtain such a compound the following 
method is adopted. Finely powdered bone-ash, containing 
normal calcium phosphate, is warmed with enough dilute 
sulphuric acid to withdraw two-thirds of the calcium, which 
unites with the sulphuric acid forming insoluble sulphate, while 
the remaining third of the calcium remains combined with 
the whole of the phosphoric acid as a soluble acid phosphate: 


H 4 Ca(P 0 4 ) 2 , or 
lime :— 


PO | (OH) 
PO) 0 2 Ca 


4 , commonly called superphosphate of 


PO | OtsCa 3 + 2S0 2 (0H) 2 = 2S0 2 -0 2 Ca + po} ( 0 2 Ca 

The clear solution is next drawn oft and concentrated in leaden 
pans, and the concentrated solution so obtained is then mixed 
with powdered charcoal, dried, and heated to low redness. During 
this process the acid phosphate loses water and becomes converted 

PO ) 

into calcium metaphosphate : Ca(P 0 3 ) 2 , or pQ 2 10 2 Ca, a com¬ 
pound which contains more phosphoric acid than the normal 
phosphate, and from which two-thirds of the phosphorus is ob¬ 
tained on-glowing with charcoal :— 

PO ) 

3 (P0 2 ) 2 0 2 Ca + 10C = P 4 + po| 0 6 Ca 3 + 10CO, 

the calcium being again converted into the normal phosphate. 
The reduction is carried on in earthenware retorts connected with 
earthenware vessels containing water. The gaseous phosphorus 
distils over and condenses under the water, while the carbonic 
oxide escapes. The raw product is then purified by redistillation 
from iron retorts and by squeezing the liquid through leather 
under warm water. Phosphorus as it occurs in trade is usually 


21 o Text-Book of Inorganic Chemistry. 

in the form of sticks ; this form is given it by the highly dangerous 
practice of sucking up the molten substance under water into 
conical glass tubes, from which it is easily removed on solidifi¬ 
cation. 

If a stick of phosphorus, one end of which is wrapped in paper, 
is used to write on a board, the characters are luminous in the 
dark. Small pieces of phosphorus are by this means rubbed oft 
on the board and gradually oxidize to phosphorous acid. The 
luminosity of phosphorus or its oxidation to phosphorous acid is 
always accompanied with the production of ozone. Various bodies 
in the gaseous state when mixed with atmospheric air prevent its 
slow oxidation and at the same time its luminosity; among these 
are small quantities of oil of turpentine vapour, ether, ammonia, 
and various hydrocarbons. Phosphorus is, however, much less 
luminous in pure oxygen than in common air, or in oxygen diluted 
with an indifferent gas. 

Phosphorus unites as easily with the halogens as with oxygen, 
liberating a large amount of light and heat. With chlorine it forms 
either phosphorus trichloride or pentachloride according to the 
relative quantities of phosphorus and chlorine which act upon one 
another. 

The affinity of phosphorus for sulphur is so great that by drop¬ 
ping flowers of sulphur on to phosphorus melted under water the 
union takes place with a very considerable elevation of temperature, 
and dangerous explosions may often be produced. Phosphorus 
does not appear to unite directly with hydrogen, nitrogen, or 
carbon. 

The principal use of phosphorus is in the manufacture of lucifer 
matches, for which both modifications are employed. Common 
matches are dipped in a mixture containing common phosphorus, 
and ignite when rubbed on any rough surface, the ignition being 
imparted to the wood by previously dipping the matches fin melted 
paraffin. In the other kind of matches, commonly known as 
safety matches , the matches themselves contain no phosphorus but 
only some substance which readily gives off oxygen (usually potas¬ 
sium chlorate), while one side of the box is covered with finely 
divided amorphous phosphorus put on with gum water. The 
matches are then ignited by rubbing them on the prepared surface. 
Their ignition is due to the change of the part of the amorphous 
phosphorus rubbed by the match into the common form by the 
heat of friction ; this then adheres to the match and causes it to 


Compounds of Phosphorus and Hydrogen . 211 

catch fire. If phosphorus is rubbed up with fat a mixture is pro¬ 
duced which is luminous in the dark, and which is used as a poison 
for rats and mice. 


COMPOUNDS OF PHOSPHORUS AND HYDROGEN. 

Two compounds of these elements are known—viz. a gaseous 
compound having the composition PH 3 and corresponding to 
ammonia (NH 3 ), and a liquid, P 0 H 4 , corresponding to amidogen 
(NH 2 ), 


Caseous Phosphoretted Hydrogren or Phosphine : PH 3 . 

This compound is a colourless gas, with an exceedingly un¬ 
pleasant garlic odour, and is nearly insoluble in water. Its specific 
gravity is 1*18, which corresponds to a molecular weight of 
1*18 + 28 , 88 = 34*08, and its molecule is therefore represented by 
the formula PH 3 . 

The gas is not spontaneously inflammable, but at once ignites 
when brought into contact with a flame, burning to form water and 
phosphoric acid. 

Phosphine is easily obtained by heating phosphorus with a 
solution of a strong base, such as potash, soda, lime, or baryta 
Although phosphorus possesses so strong an attraction for oxygen 
it cannot decompose water even at the boiling temperature, but the 
decomposition goes on readily in the presence of a strong soluble 
base, a portion of the phosphorus uniting with the base to form a 
hypophosphite. The gas may also be prepared from solid calcium 
phosphide, which, in contact with warm water, is rapidly decom¬ 
posed into phosphine, calcium hypophosphite, and calcium hydrate. 
But the gas obtained by either of these processes is always spon¬ 
taneously inflammable. This is due to the fact that it contains 
small quantities of the volatile liquid phosphoretted hydrogen, 
which is spontaneously inflammable, and causes the gas to ignite 
in the air just as it does when mixed with other inflammable gases, 
such as carbonic oxide or hydrogen. The following equations re¬ 
present the simultaneous production of the two compounds when 
phosphorus is heated with caustic potash :— 



212 Text-Book of Inorganic Chemistry. 

(i.) 4P + 3KOH + 3H 2 0 = 3H 2 PO-OK + PH 3 . 

Potassium hypo- 
phosphite 

(ii.) 6P + 4KOH + 4H 2 0 = 4 H 2 PO-OK + P 2 H, 

The preparation of phosphine containing the spontaneously 
inflammable liquid compound must be performed with care. If 
phosphorus were heated in a flask half filled with caustic potash 
and provided with a tightly fitting cork and delivery-tube, the 
phosphoretted hydrogen would unite with the oxygen in the flask 
and catch fire, and the heat so produced would be sufficient to 
cause an explosion. It is better, therefore, to fill the flask nearly 
up to the neck and to place the cork in loosely until the oxygen 
contained in the small quantity of air present has been consumed. 
The delivery-tube must be wide so that it may not become closed 
with small particles of phosphorus carried over mechanically ; it 
should dip under warm water. 

In order to prepare the spontaneously inflammable gas from 
calcium phosphide, this compound is broken into small pieces of 
the size of a pea and dropped through a wide tube into a Woulffe’s 
bottle half filled with warm water, the other tubulus of the bottle 
being provided with a delivery tube. The wide tube must just dip 
under the liquid and must be closed with a small cork. Before any 
calcium phosphide is introduced, the air contained in the Woulffe’s 
bottle must be completely displaced by carbonic acid to prevent 
the gas from catching fire inside the bottle. 

The spontaneously inflammable gas when preserved over water 
for some time, and particularly when exposed to the light, is found 
to lose its property of catching fire when brought into contact with 
the air, and at the same time the walls of the glass vessel become 
covered with a yellow film. 1 This phenomenon depends upon the 
instability of the liquid compound, which easily breaks up, espe¬ 
cially in sunlight, into phosphine and phosphorus :— 

3 P 2 H 4 = 4PH3 + 2P. 

The spontaneously inflammable liquid may be easily separated 
from the gaseous compound by passing the mixture through a 
U-tube placed in a freezing mixture. By this means the volatile 
liquid is condensed and the gas is no longer spontaneously inflam¬ 
mable. 

1 This yellow film is considered by some chemists to contain hydrogen and 
to consist of a compound of 4 atoms of phosphorus united with 2 atoms of 
hydrogen : P4H2 .—Ed. 


Compounds of Phosphorus and Oxygen. 213 

Pure phosphine may be prepared by dropping calcium phos¬ 
phide into hydrochloric acid instead of into water, or by acting 
tipon phosphorus with alcoholic instead of aqueous caustic potash. 
Hydrochloric acid and alcohol appear to prevent the formation of 
the liquid phosphoretted hydrogen, or to decompose it immedi- 

i; atel y- 

The pure gas when led into nitric acid mixed with a small 
quantity of the fuming acid becomes spontaneously inflammable. 
This is probably due to the abstraction of a portion of its hydrogen> 
which is oxidized by the nitric acid, with the production of liquid 
I phosphoretted hydrogen. 

Phosphine is distinguished from its analogue ammonia not only 
by its insolubility in water, but also by its incapability of uniting 
either with hydrochloric acid or with oxy-acids to form chemical 
compounds, but it still possesses weak basic properties. If dry 
phosphine and dry hydriodic acid are brought into contact with one 
i another they unite to form a compound called phosplionium iodide , 
PH 4 I (p. 128), analogous in composition to ammonium iodide, 
which crystallizes in cubes, and may be sublimed unchanged. The 
compound is, however, unstable ; it is decomposed by water into 
hydriodic acid and phosphine. 


liquid Phosphoretted Hydrogen : P 2 H 4 . 

This compound, which is prepared from the spontaneously in¬ 
flammable gas by passing it through a freezing mixture, is a colour¬ 
less, powerfully refractive liquid, does not mix with water, and at 
once catches fire when exposed to the air. The liquid undergoes a 
gradual decomposition, even in sealed tubes, into phosphine and 
amorphous phosphorus. The gas finally exercises so great a 
pressure on the tube as to produce an explosion. It is, therefore, 
not practicable to preserve the liquid in sealed tubes. 


COMPOUNDS OF PHOSPHORUS AND OXYGEN. 

Although phosphorus is so closely allied to nitrogen, we are 
only acquainted with two oxides of this element—viz. phosphorous 
anhydride (P 2 0 3 ) and phosphoric anhydride (P s 0 6 ), corresponding 
to nitrous anhydride and nitric anhydride respectively. Compounds 






214 


Text-Book of Inorganic Chemistry. 

of phosphorus having a similar composition to nitric peroxide, 
nitric oxide, and nitrous oxide, have not yet been prepared. 

From these two oxygen compounds are derived the two acids, 
common tribasic phosphoric acid and dibasic phosphorous acid. 
Besides these, a third (lower) acid is also known, which is called 
hypophosphorous acid and is monobasic. The following are the 
formulae of these three acids of phosphorus, all containing pentad 
phosphorus :— 

(OH 

Common phosphoric acid . . . PCU OH 

[OH 

Phosphorous acid .... HPo|q£| 
Hypophosphorous acid . . . H 2 PO-OH 

A fourth acid of phosphorus has recently been discovered of 
the composition : jpo(OH) 2 ’ and called hypophosphoric acid. 


PHOSPHORIC ANHYDRIDE : P 2 0 3 . 

The white fumes produced when phosphorus burns with a flame 
in dry air consist of this compound. If the phosphorus is burnt in 
a closed space— eg. under a bell-jar—they then condense to a light 
snow-white powder, which very readily attracts moisture from the 
air and forms liquid phosphoric acid. When thrown into water 
the anhydride hisses like a red-hot iron, and a large quantity of 
heat is at the same time set free. In fact, there are few substances 
which equal this body in its attraction for water. Unlike nitric 
anhydride, it is not volatile at ordinary temperatures, and does not, 
therefore, fume in moist air ; but it may be sublimed when strongly 
heated in a glass tube. 

The acid first produced when the anhydride unites with water 
contains only one atom of hydrogen displaceable by a metal, and is 
therefore monobasic. This acid is metafihosphoric acid\ P 0 2 - 0 H, 
and its production is shown by the following equation :— 

PO ^ 

PQ^ j f-) + H 2 0 = 2 P 0 2 - 0 H. 

Phosphoric Metaphosphoric 

anhydride acid 



215 


Phosphoric Anhydride. 


Besides this phosphoric acid, which corresponds to nitric acid in its 
composition, we are also acquainted with a tri- and a tetrabasic acid. 
Tribasic phosphoric acid, which has the composition PO(HO) 3 , 
and whose normal salts with monobasic metals contain three atoms 
of the metal replacing the three atoms of hydrogen, is common or 
orthophosphoric acid. Only compounds of this tribasic acid occur 
in nature. 

Metaphosphoric acid, produced as we have seen from phos¬ 
phoric anhydride and water, takes up a molecule of water when 
boiled and becomes converted into common phosphoric acid 


P 0 2 - 0 H + H 2 0 - PO(OH) 3 . 

The change which has taken place may be readily proved by a few 
simple reactions. Metaphosphoric acid coagulates white of egg, 
common phosphoric acid does not; on adding a little ammonia, 
the former gives with silver nitrate a white precipitate of silver 
metaphosphate : P 0 2 - 0 Ag, the latter with the same reagent gives 
a yellow precipitate of silver orthophosphate : PO(OAg) 3 . The 
conversion of metaphosphoric acid into common phosphoric acid 
is therefore complete when the liquid no longer coagulates white of 
egg and gives a bright yellow precipitate with silver nitrate. 

Preceding the formation of common phosphoric acid from meta¬ 
phosphoric acid, when the solution is warmed, is the production of 
tetrabasic pyrophosphoric acid, which may be recognized by the 
fact that this acid does not coagulate white of egg but gives a white 
precipitate of silver pyrophosphate when mixed with silver nitrate. 
Its production from the monobasic acid is shown by the following 


equation:— 


2P0 2 -OH + H 2 0 

Metaphosphoric 
acid 


^JPOtOH), 

U |PO(OH) ? 

Pyrophosphoric 
acid 


All three acids may therefore be represented as compounds of 
phosphoric anhydride with varying quantities of water, thus — 

po 2 1 0 

po 2 ) u 


h 2 o 


= 2P0 2 OH 


PO, 

po: 


no 


2 H 2 0 


Metaphosphoric 

acid 

]PO(OH) 2 

u 1po(oh) 2 

Pyrophosphoric 


p8:}° + 3 H.O - 


2 PO(OH) 3 

Orthophosphoric 

acid 




216 


Text-Book of Inorganic Chemistry . 


ORTHO- OR COMMON PHOSPHORIC ACID : PO(OH) 3 . 

On evaporating the aqueous solution of common phosphoric acid, 
when the temperature is not allowed to rise much above 150°, the 
acid remains behind as a colourless, thick, viscid liquid, of specific 
gravity i *88, from which large, colourless crystals occasionally 
separate out. It readily attracts moisture from the air and 
deliquesces. When heated more strongly it loses water and 
becomes first converted into pyro- and then into metaphosphoric 
add. 

Common phosphoric acid is easily prepared by oxidizing 
phosphorus with nitric acid. Sticks of phosphorus are heated in a 
retort with pure nitric acid of specific gravity 1*2, when an energetic 
reaction takes place and red vapours of the lower oxides of 
nitrogen are given off. The retort is best provided with a cooled 
receiver, so that the nitric acid distilling over may be poured back 
into the retort. As soon as all the phosphorus has disappeared 
the excess of nitric acid is driven off, which at the same time 
oxidizes the phosphorous acid, also present, to the higher phosphoric 
acid. Finally, the viscid liquid is heated in a platinum dish until 
the temperature reaches 150°. 

Another and very easy method of preparing phosphoric acid is 
from phosphorus pentachloride. If this substance is dropped into 
water in small quantities at a time, it decomposes, with consider¬ 
able rise in temperature, into phosphoric and hydrochloric acids, 
phosphoric oxychloride being produced as an intermediate pro¬ 
duct : — 

PC 1 5 + 4H 2 0 = PO(OH) 3 + 5HCI. 

Pure phosphoric acid may then be obtained by evaporating 
the strongly acid liquid. 

Since commercial phosphorus always contains arsenic, because 
the sulphuric acid used in the preparation of the acid calcium 
phosphate is always contaminated with arsenic, the phosphoric 
acid prepared either from phosphorus or from phosphorus penta¬ 
chloride always contains traces of arsenic in the form of arsenic 
acid. In order to separate this impurity the phosphoric acid, after 
being freed from the nitric or hydrochloric acid, is diluted with a 
large quantity of water, gently warmed, and a stream of sulphu¬ 
retted hydrogen passed through it as long as insoluble yellow 
arsenious sulphide separates out. The whole of the arsenic is. 


Ortho- or Common Phosphoric Acid. 


217 


precipitated, when a small portion of the clear liquid again 
saturated with sulphuretted hydrogen and allowed to stand for 
some hours shows no further turbidity. The clear filtered liquid 
then yields pure phosphoric acid on evaporation. 

Phosphoric acid may also be prepared from the acid calcium 
phosphate (superphosphate of lime), obtained by acting on bone-ash 
or some other form of calcium phosphate with sulphuric acid. 
The solution of the superphosphate, which still, contains calcium 
sulphate, is mixed with a solution of ammonium carbonate to pre¬ 
cipitate the lime. The filtered liquid, which contains ammonium 
sulphate and phosphate, is then evaporated to dryness and glowed 
in order to drive off the ammonium sulphate and to expel the 
ammonia from the ammonium phosphate. Phosphoric acid pre¬ 
pared in this way always contains, however, small quantities of 
ammonium salts. 

Phosphoric acid is one of the strongest inorganic acids. At a 
high temperature it even expels sulphuric acid from its compounds, 
being less volatile than this acid. Since it is a tribasic acid, it forms 
three series of salts :— 

(i.) Normal salts, in which all three of its hydrogen atoms are 
displaced by a metal. 

(ii.) Monacid salts , in which two atoms of hydrogen are dis¬ 
placed by a metal, and (iii.) Diacid salts , in which only one atom of 
hydrogen is displaced. 

In the normal salts the hydrogen may be either displaced by 
three atoms of a monad metal, by one of a monad and one of a 
dyad, or by one of triad metal, thus :— 

Normal sodium phosphate 

Normal sodium-ammonium phosphate 


Normal ammonium-magnesium phosphate 


PO(ONa) 3 

PO i (ONa) 2 
fU (ONH 4 

rU |ONH 4 

PO-OgFe'" 


Normal ferric phosphate ... 

In the first class of acid salts the two atoms of hydrogen may 
either be displaced by two atoms of a monad metal or by one atom 
of a dyad, thus :— 

Monacid sodium phosphate 


PO I 
FU OH 


Monacid calcium phosphate 


PO 


0 2 Ca' 

OH 


218 Text-Book of Inorganic Chemistry. 


And, finally, for the second class of acid salts we may take as an 
example:— 


Diacid sodium phosphate 


PO 


(ONa 
I (OH), 


The phosphates of the metals, particularly the normal salts, 
are mostly insoluble in water, but easily soluble in nitric acid ; only 
the phosphates of the alkalies can be obtained in a well-defined 
crystalline state from their solutions. 

Normal silver phosphate is a bright yellow amorphous precipi¬ 
tate, easily soluble in nitric acid. It is obtained on the addition of 
silver nitrate to a solution of common sodium phosphate. This 
latter compound, the phosphate of soda of the shops, is not the 
normal compound, but is monacid sodium phosphate— i.e. it con¬ 
tains one atom of hydrogen undisplaced. The reaction by which 
normal silver phosphate is produced when this salt is mixed with 
silver nitrate is as follows :— 

POjg^* 3N0 2 -OAg = P 0 ( 0 Ag) 3 + 2N0 2 -0Na + N 0 2 0 H. 

Whence it is seen that nitric acid is at the same time set free, and 
if the clear liquid is filtered off, it will be found to possess an acid 
reaction, and to contain some silver phosphate dissolved in it. 
We produce, therefore, from the two neutral solutions of mono- 
liydric sodium phosphate and silver nitrate an acid liquid. The 
phosphoric acid contained in this and other similar acid compounds 
can, therefore, only be completely precipitated by silver nitrate by 
neutralizing with soda or some other suitable base. 

The insolubility of ammonium-magnesium phosphate in water, 
especially in presence of ammonia, and the constant composition of 
this compound, makes it well adapted both to detect phosphoric 
acid and to determine its quantity. It is precipitated as a white 
crystalline powder when a solution of sodium or any other soluble 
phosphate is mixed with magnesium sulphate, ammonium chlo¬ 
ride, and ammonia. 

If monacid sodium phosphate is mixed with magnesium sul¬ 
phate the acid magnesium compound is produced. Thus :— 

po |L°H a)a + S0 2 .0 2 Mg = P0{g k Mg + s<V(ONa) a , 

which is soluble in water, even in the presence of ammonium chloride. 
If, however, ammonia is added in excess, the insoluble ammonium- 
magnesium phosphate is precipitated :— 


219 


Pyrophosphoric Acid. 



o 2 m g 

OH 


+ NH, = 


PO 


( 0 .,Mg 
(ONH 4 


The addition of ammonium chloride is necessary to prevent the 
precipitation of magnesium hydrate from the excess of magnesium 
sulphate and ammonia. Magnesium hydrate forms an easily 
soluble compound with ammonium chloride. 

An excellent method of detecting and separating phosphoric 
acid in nitric acid solutions containing iron, lime, &c., is furnished 
by molybdic acid. If such a solution containing free nitric acid is 
mixed with a solution of ammonium molybdate in nitric acid, the 
liquid first becomes yellow and then deposits a yellow crystalline 
precipitate consisting of a compound of phosphoric and molybdic 
acids. By this means a soil, for example, may be examined for 
phosphoric acid, and its quantity, if present, determined. 

If solid monacid sodium phosphate is gently heated, it first of 
all loses its water of crystallization. On further heating it gives up 
the hydrogen which it contains, together with sufficient oxygen to 
form water, and normal sodium pyrophosphate remains behind :— 


2PO 


((ONa) 2 
{OH 


(PO(ONa) 2 
jPO(ONa) 


H 2 0 . 


When diacid sodium phosphate is heated, water is also given 
off, but the compound which remains behind is then normal 
sodium metaphosphate :— 

PO !(OH) 2 '= POj'ONa + H 2 0. 

The same compound is more easily obtained by heating acid 
ammonium-sodium phosphate (microcosmic salt), which then gives 
off ammonia and water :— 


fONa 

PO ONH 4 - P 0 2 - 0 Na + NH 3 + H 2 0 . 

(OH 


Pyrophosphoric Acid 


• oi P °(° H >3- H P O 

• u 1po(oh) 2 ~ n * r *y r 


This compound is obtained as a colourless crystalline mass, 
easily soluble in water, when phosphoric acid is heated slightly 
above 200°, until a portion when neutralized with ammonia gives a 
pure white precipitate with silver nitrate. If sodium pyrophos¬ 
phate, obtained as mentioned above, is dissolved in water mixed 





220 


Text-Book of Inorganic Chemistry. 


with silver nitrate, the white precipitate filtered off, washed well, and 
then again decomposed by passing a current of sulphuretted hydro¬ 
gen through water containing it in suspension, black silver sulphide 
is formed, and the acid solution contains pyrophosphoric acid. 
Such an aqueous solution of the acid cannot, however, be kept for 
any time ; it combines with water and forms common phosphoric 
acid, slowly when cold, more rapidly when heated :— 


O 


(PO(OH) 2 
|po(OH) 2 


+ h 2 0 


2 PO(OH) 3 . 


Pyrophosphoric acid is a tetrabasic acid, its four atoms of 
hydrogen may be displaced by the same or by different metals, and 
besides the normal salts three classes of acid salts are also known 
accordingly as one, two, or three atoms of hydrogen are displaced. 


Rftetaphosplioric Acid : P 0 2 *OH. 

This acid is obtained in aqueous solution by allowing phos¬ 
phoric anhydride to deliquesce in the air, or as a vitreous, 
amorphous mass by heating common phosphoric acid to redness. 
In the latter state it forms the glacial phosphoric acid of the 
pharmacists. In combination with sodium it is produced, as we 
have already seen, by heating microcosmic salt to redness. 

The acid and its soluble salts give with silver nitrate an 
insoluble white precipitate of silver metaphosphate: P 0 2 * 0 Ag, 
and are distinguished from the two other phosphoric acids and 
their salts by their property of coagulating albumen (white of egg). 

Metaphosphoric acid when in solution gradually absorbs water 
and is converted into common phosphoric acid. The change goes 
on slowly in the cold, but more rapidly when the liquid is heated. 
Solutions of the metaphosphates also take up water in the same 
way— e.g .:— 

PCVONa + H 2 0 = P °{( 0 H) a - 

The recently discovered hypophosphoric acid : j pq^qH) 2 ’ 1S a 

tetrabasic acid intermediate between phosphoric and phosphorous 
acids, and is produced with phosphorous and phosphoric acids, 
when phosphorus oxidizes slowly in moist air. Its solution may be 
evaporated at a low temperature, but is decomposed on boiling into 
phosphorous and phosphoric acids. 


Phosphorous A nhydride. 


221 


Phosphorous Anhydride : P 2 0 3 . 

The white fumes produced when a stick of phosphorus is 
exposed to the air consist principally of phosphorous acid. The 
pure anhydride may be obtained by gently heating a stick of dry 
phosphorus in a glass tube and leading a current of dry air over it. 
The phosphorus then burns with a pale greenish flame, and the 
anhydride is deposited in the cool end of the tube as a snow-white 
powder resembling phosphoric anhydride. Phosphorous anhydride 
has still a strong attraction for oxygen ; when exposed to the air it 
catches fire and burns to phosphoric anhydride. 

When phosphorous anhydride is brought into contact with 
water, or when phosphorus is allowed to oxidize slowly in moist 
air, the corresponding acid—phosphorous acid—is produced. 

Phosphorous Acid : HPo|q^ 

Although this compound contains three atoms of hydrogen, 
only two of them can be displaced by metal, and it is therefore a 
dibasic acid. It is best prepared by dropping phosphorus trichlo¬ 
ride into water. The chloride then breaks into hydrochloric and 
phosphorous acids, thus :— 

PC 1 3 + 3H 2 0 = 3HCI + HP0{gH 

The strongly acid liquid is freed from the hydrochloric acid by 
rapid evaporation, when phosphorous acid remains behind as a 
clear, colourless syrup, which if preserved for some time over sul¬ 
phuric acid in a vacuum changes into a crystalline mass. 

Pure phosphorous acid is obtained as a bye-product in the pre¬ 
paration of acetyl chloride, when glacial acetic acid is gradually 
added to phosphorus trichloride, contained in a retort, and the 
volatile acetyl chloride together with the excess of acetic acid dis¬ 
tilled off. Phosphorous acid then remains behind as a thick 
colourless liquid, which solidifies on cooling. If we represent the 
radical of acetic acid by the symbol A, acetic acid then has the 
formula AOH, and the process is represented by the following 
equation :— 

PC 1 3 + 3AOH - 3ACI + HPO(OH) 2 . 

Phosphorous acid melts at 70°; if the acid is strongly heated it 





222 Text-Book of Inorganic Chemistry. 

breaks up into phosphoric acid and phosphine, the latter catching 
fire in the air :— 

4 HPO(OH ) 3 = 3 PO -(OH ) 3 + ph 3 . 

The phosphites when strongly heated undergo a similar change. 

The strong attraction of phosphorous acid for oxygen and its 
readiness to pass into phosphoric makes it a powerful reducing 
agent. When an aqueous solution is gently warmed with solutions 
of mercury or silver the compounds of these metals are reduced to 
the metallic state. Similarly, when the acid is exposed to the air it 
gradually takes up oxygen and becomes converted into phosphoric 
acid. 


Hypophosphorous Acid : (HoPO)OH. 

This acid of phosphorus, which contains less oxygen than phos¬ 
phorous acid, is monobasic. Its molecule contains three atoms 
of hydrogen, but only one of these atoms can be displaced by 
a metal. The acid is soluble in water, and on evaporating down 
its aqueous solution remains behind as a syrupy liquid, which crys¬ 
tallizes below o°. Most of its salts crystallize easily, and are all 
soluble in water. 

The hypophosphites are produced, as we have already seen 
(p. 211), when an aqueous solution of an alkali is heated with phos¬ 
phorus. Phosphoretted hydrogen is then evolved, and the hypo- 
phosphite remains behind in solution. For the preparation of 
hypophosphorous acid and its salts it is best to employ barium 
hydrate (baryta-water), since the insolubility of barium carbonate 
affords a ready means of removing the excess of alkali. After a 
hot saturated solution of baryta-water has been heated with excess 
of phosphorus for some time until phosphoretted hydrogen is no 
longer evolved, the whole is allowed to cool, and carbonic acid 
passed through the clear liquid until it no longer reacts alkaline ; 
heat being applied towards the end. The barium carbonate is 
then filtered off, and the liquid evaporated down, when barium 
hypophosphite crystallizes out on cooling. The production of this 
salt is represented by the following equation, which is similar to 
that by which potassium hypophosphite is formed (p. 212) :— 

6 P + 2Ba(OH) 2 + 4 H s O = 2^o}°’ Ba " + P 2 H 4 . 



Compounds of Phosphorus with the Halogens. 223 

Milk of lime (calcium hydrate) may be employed instead of the 
baryta-water, when calcium hypophosphite is easily obtained in the 
crystalline form. Calcium hypophosphite is also obtained when 
calcium phosphide is decomposed by warm water (p. 211). 

The free acid may be prepared from the barium salt by pre¬ 
cipitating the barium with sulphuric add and filtering off from the 
insoluble barium sulphate. The aqueous solution can be evapo¬ 
rated to a certain degree of concentration without decomposition, 
but when more strongly heated breaks up (like phosphorous acid) 
into phosphoric acid and phosphine :— 

2 H 2 PO-OH = PO(OH) 3 + PH 3 . 

It possesses a strong attraction for oxygen, and, like phosphorous 
acid, reduces various metals from their oxides or from solutions of 
their salts. 


COMPOUNDS OF PHOSPHORUS WITH THE 
HALOGENS. 

We are acquainted with two compounds of phosphorus and 
chlorine—viz. phosphorus trichloride, PC 1 3 , in which the phos¬ 
phorus is triad, and phosphorus pentachloride, PC 1 5 , in which the 
phosphorus is pentad. Similar compounds with bromine are also 
known. A corresponding phosphorus triodide, PI 3 , also exists, but 
the pentaiodide has not been prepared; on the other hand, a com¬ 
pound containing less iodine, of the composition (PI 2 ) 2 or P 2 I 4 — 
phosphorus di-iodide—is known. Phosphorus forms only one com¬ 
pound with fluorine—the gaseous pentafluoride, PF 5 . 

The attraction between phosphorus and the halogens is so great 
that all these compounds, except the fluoride, can be easily pre¬ 
pared by the direct union of their constituents : while the strong 
attraction of phosphorus for oxygen and of the halogens for hydro¬ 
gen is the reason why these compounds are easily decomposed by 
water with a considerable evolution of heat. 





224 


Text-Book of Inorganic Chemistry . 


PHOSPHORUS TRICHLORIDE (Phosphorous Chloride), j 
Composition : PC 1 3 . 

This compound is a colourless liquid, fuming in moist air, and 
which may be prepared in quantity in the following manner. The 
bottom of a large tubulated retort is covered about half an inch 
deep with dry sand, and when the air has been expelled by car¬ 
bonic acid carefully dried sticks of phosphorus are introduced. 
The neck of the retort is then connected with a cooled receiver, 
and a rapid stream of dry chlorine led into it. The phosphorus 
catches fire in the atmosphere of chlorine, melts, and burns with a 
pale flame, forming the trichloride which distils over. The process 
must be so regulated that the chlorine is not in excess, partly by 
reducing the current of the gas and partly by heating the retort 
more strongly. An excess of chlorine would produce the solid 
pentachloride. But if too much phosphorus vapour is produced, 
a portion distils over into the condenser and receiver. In the pre¬ 
paration of large quantities it is best to employ several vessels 
generating chlorine, when the trichloride can then be distilled in an 
even stream. 

The substance so obtained is never pure ; either it contains 
dissolved pentachloride or free suspended phosphorus. In the 
latter case redistillation is sufficient to remove most of the phos¬ 
phorus : in the former case the liquid must be digested with free 
phosphorus, which converts the pentachloride into the trichloride, 
and be again distilled. The phosphorus trichloride then only con¬ 
tains small quantities of free dissolved phosphorus, for which it is 
a good solvent. 

Phosphorus trichloride is a clear, colourless liquid of 1*58 
specific gravity. It boils at 76°, and its vapour has a density of 
47, which corresponds to a molecular weight of 137, and its com¬ 
position PC 1 3 . It fumes in moist air and is at once decomposed 
by water into phosphorous and hydrochloric acids with a consider¬ 
able rise in temperature. With chlorine it unites readily and forms 
the solid pentachloride. 



Phosphorus Pentachloride (.Phosphoric Chloride). 225 


PHOSPHORUS PENTACHLORIDE (Phosphoric 
Chloride). 

Composition : PC1 5 . 

The solid volatile phosphorus pentachloride is easily prepared 
from the liquid trichloride by causing the latter to unite directly 
with two further atoms of chlorine. A wide-necked flask is about 
half filled with phosphorus trichloride on to which chlorine is led 
by a wide tube passing through a glass plate closing the neck of 
the flask. The tube must not touch the liquid or it will become 
closed by the solid pentachloride. The chlorine is rapidly ab¬ 
sorbed, a large quantity of heat is evolved, and the flask must be 
kept cool throughout the operation. It is necessary to stir the 
mass from time to time and finally to slightly warm it. The con¬ 
version is complete when chlorine is no longer absorbed on allow¬ 
ing the solid mass to stand in contact with the gas in closed 
vessels. 

Another method of preparation consists in dissolving phos¬ 
phorus in carbon disulphide, and then passing chlorine through 
this solution. Phosphorus pentachloride, which is insoluble in 
carbon disulphide, then separates out in indistinct crystals. 

Phosphorus pentachloride is a white crystalline mass fuming 
in the air. It sublimes without melting at ioo°, and melts under 
pressure at about 148. 0 Its vapour when inhaled is exceedingly 
irritating, a fact which should not be forgotten in experimenting 
with it. All operations with it should be carried on in a good 
draught of air. Its vapour density is 3-6, corresponding to a 
molecular weight of 104, which is only one-half that required for 
the formula PCI- (208*5). The reason of this is that one molecule 
of the pentachloride breaks up when heated (dissociates) into one 
molecule of the trichloride and one of free chlorine, which occupy 
twice the volume that one molecule of the pentachloride would 
take up. On cooling again the two substances recombine and 
produce phosphorus pentachloride. When the pentachloride is 
thrown into water it is at once decomposed with a large evolution 
of heat, while a colourless oil collects at the bottom of the vessel 
and after a time disappears. This oil is phosphoric oxychloride, 

Q 


226 Text-Book of Inorganic Chemistry 

P0C1 3 , or the chloride of common phosphoric acid, PO(OH) s . It 
is decomposed by water into phosphoric and hydrochloric acids: — 

PC1 5 + H 2 0 = POCI3 4 2HCI. 

POCI3 + 3H.O = PO-(OH ) 3 4- 3HCI. 

The tendency of phosphorus pentachloride to exchange two 
atoms of chlorine for one of oxygen makes it a valuable agent for 
preparing compounds of chlorine which are difficult to obtain in 
any other way. We have seen (p. 165) that when a mixture of 
phosphorus pentachloride and sulphuric acid is heated, chlorsul- 
phonic acid, phosphoric oxychloride, and hydrochloric acid are 
formed. 


PHOSPHORIC OXYCHLORIDE (Pbosphoryl 
Chloride). 


Composition : POCl 3 . 


This substance, which closely resembles phosphorus trichloride 
in its physical properties, is gradually produced from phosphorus 
pentachloride when exposed to moist air. It is, however, better 
prepared by heating the pentachloride with certain acids or with 
compounds containing water chemically combined. The best 
substance to employ for this purpose is ordinary crystallized boric 
acid, which is then converted into boric anhydride :— 


3PC1, 


2B(OH) s 

Boric acid 


- 3POCI3 


4 B 2 0 3 4 

Boric anhydride 


6HC1. 


Phosphorus pentachloride and boric acid are mixed in the 
proportions shown in the equation—*>. about five parts of the former 
with one part of the latter—the mixture placed in a retort, and then 
heated. Torrents of hydrochloric acid are given off, and phosphoric 
oxychloride distils over into the well-cooled receiver. The boric 
anhydride, which finally remains behind in the retort, may be 
dissolved up in warm water, allowed to crystallize, and again used 


for the same purpose. 

Phosphoric oxychloride is a colourless liquid sinking in water, 
by which it is soon decomposed. It has a specific gravity of 17, 
boils at iio°, fumes when exposed to the air, and possesses a 
piercing odour. It is easily decomposed by water into hydrochloric 
and phosphoric acids. 

Corresponding to phosphoric oxychloride is a sulphur compound, 
phosphoric sulphochloride : PSC1 3 , which is also a colourless liquid, 




Bromides of Phosphorus. 227 

of specific gravity r6 and boiling at 125° It fumes in the air and 
is slowly decomposed by water into phosphoric acid, hydrochloric 
acid, and sulphuretted hydrogen. This compound is obtained by 
heating phosphorus pentachloride with certain metallic sulphides— 
e.g. antimonous sulphide :— 

3 PC 1 5 + Sb 2 S 3 = 3 PSC 1 3 + 2 SbCl 3 . 

Antimonous Antimonous 

sulphide chloride 


BROMIDES OF PHOSPHORUS. 

Phosphorus Tribromide (Phosphorous Bro?nide) : PBr 3 . 

This is a heavy colourless liquid, of specific gravity 2*9, boiling 
at i 73°j an d closely similar in its chemical properties to the 
corresponding trichloride. It is best obtained by dissolving a 
known weight of phosphorus in about eight times as much carbon 
disulphide, and then dropping in the requisite quantity of bromine— 
i.e. three atoms of bromine for every atom of phosphorus. As soon 
as the reaction has ceased and all bromine has disappeared the 
carbon disulphide is evaporated off on a water bath, and then the 
phosphorus tribromide distilled over. 


Phosphorus Pentabromide (Phosphoric Bromide) : PBr 5 . 

The pentabromide is prepared in precisely the same way as 
the tribromide, except that five atoms of bromine for every atom 
of phosphorus are employed. 

It is a yellow, crystalline solid, less stable than phosphorus 
pentachloride, decomposing even at ioo° into phosphorus tribromide 
and bromine. When acted upon by water it behaves in a similar 
manner to the corresponding pentachloride. 


Phosphoric Oxybromide : POBr 3 , is a crystalline solid, melting 
at 45 0 and boiling at 195° It is obtained, together with acetyl 
bromide and hydrobromic acid, by acting on phosphorus penta¬ 
bromide with acetic acid :— 

PBr 5 + AOH = POBr 3 + ABr + HBr. 




228 


Text-Book of Inorganic Chemistry. 


IODIDES OF PHOSPHORUS. 

We have already remarked that phosphorus penta-iodide is 
unknown ; only the tri- and di-iodides have been as yet prepared. 
Both are obtained in the same way—viz. adding the requisite 
quantity of iodine to a solution of phosphorus in carbon disulphide. 


Phosphorus Tri-iodide : PI 3 , is soluble in carbon disulphide, 
but separates out as red crystals when the liquid, after distilling off 
the disulphide, is placed in a freezing mixture. It melts at 55°, 
gives off iodine at its boiling temperature, and is easily decomposed 
by water into phosphorous and hydriodic acids. 


Phosphorus Di-iodide : P 2 I 4 , can be obtained not only by the 
method indicated above, but also by the direct union of solid phos¬ 
phorus and iodine. For this purpose a flask is filled with carbonic 
acid gas and the phosphorus and iodine introduced in the correct 
proportions. On contact they combine energetically with one 
another, forming a dark-coloured liquid, which, when the reaction 
has been completed by slight rise in temperature, solidifies to a 
crystalline mass on cooling. When this mass is dissolved in hot 
carbon disulphide, phosphorus di-iodide is deposited on cooling 
as bright orange coloured tablets or prisms. The di-iodide melts 
at no°to a bright red liquid, and is decomposed by water into 
hydriodic acid, phosphorous acid, and amorphous phosphorus :— 
3P 2 I 4 + I2H 2 0 = 4HPO(OH) 2 + 12HI + 2P. 


COMPOUNDS OF PHOSPHORUS AND SULPHUR. 

Phosphorus and sulphur unite in several proportions with one 
another. They have a strong attraction for one another, and their 
union is accompanied with so considerable an evolution of heat 
that if large quantities are employed disastrous explosions may 
be produced. 

But if amorphous phosphorus is employed, large quantities of 
these compounds may be easily and safely prepared. Powdered 



Arsenic. 


229 


sulphur and amorphous phosphorus are mixed together, and the 
mixture heated in a closed Hessian crucible. The crucible is first 
gently heated, and then more strongly after the reaction has 
taken place. The two substances unite with one another at the 
moment when the amorphous phosphorus passes into the common 
form. A large quantity of heat is given off, but no explosion takes 
place. If sulphur is in excess phosphorus pentasulphide : P 2 S 6 , 
is produced, while with excess of phosphorus tetraphosphoric 
trisulphide : P 4 S 3 , is formed. 

Phosphorus Pentasulphide {Phosphoric Sulphide ) : P 2 S 5 , 
may also be prepared by heating yellow phosphorus with sulphur 
and carbon disulphide in sealed tubes at 210°. On cooling, the 
compound separates out in long yellow needles. It melts at 210°, 
boils without decomposition at about 530°, and is decomposed by 
water into phosphoric acid and sulphuretted hydrogen. 

Tetraphosphoric Trisulphide : P 4 S 3 , is easily soluble in cold 
carbon disulphide, from which it crystallizes in long, yellow, rhombic 
prisms, which melt at 166 0 and are only slowly decomposed by 
water into phosphorous acid and sulphuretted hydrogen. 

Among other sulphides of phosphorus may be mentioned : 
Phosphorus disulphide : PS 2 , or more probably P 3 S 6 , which is also 
prepared by heating a mixture of phosphorus, sulphur, and carbon 
disulphide. 


ARSENIC. 

Chemical Symbol : As .—Atomic Weight'. 75. 

Arsenic occurs in nature both native and in chemical combina¬ 
tion with other elements. Native arsenic is sometimes found in 
the crystalline^state, but it generally occurs in rough lumps, which 
easily break up into uneven laminae. More common in nature are 
its compounds, of which the following are the most important: 
Arsenical iron, FeAs 2 ; arsenical iron pyrites ox mispickel, FeSAs ; 
Kupfer-nickel , NiAs ; smaltine or tin white cobalt (Co,Ni,Fe) As 2 ; 
realgar, As 2 S 2 ; orpiment, As 2 S 3 . Besides these, arsenic is also 
found in combination with oxygen as white arsenic : As 2 0 3 , and in 



230 Text-Book of Inorganic Chemistry . 

9 

the form of arsenic acid in various minerals, such as phenacolite , 
cobalt-bloom , mimetesite , &c. Finally, it is a very remarkable fact 
that arsenic, probably as arsenious acid, has been discovered in 
many mineral springs, in those of Ems, Kissingen, Pyrmont, 
Schwalbach for example. The quantity is, however, so extremely 
small that the physiological action of the mineral waters can 
scarcely be ascribed to its influence. 

Arsenic (commonly called ‘ metallic ’ arsenic to distinguish it 
from ‘white’ arsenic, the oxide), when freshly broken, possesses a 
steel-grey colour, and strong metallic lustre. It is very brittle, 
and may therefore easily be reduced to powder. Its specific 
gravity is 57. When heated arsenic sublimes and condenses 
partly as an amorphous mass, and partly as rhombohedra ; it may 
be melted under increased pressure. The specific gravity of its 
vapour at about 86o° is 10-2, which corresponds to a molecular 
weight of about 300, or, in other words, its molecule in the gaseous 
state contains four atoms. In this abnormal property arsenic re¬ 
sembles, therefore, phosphorus and sulphur. 

When exposed to moist air arsenic loses its metallic lustre, and 
becomes dull in consequence of surface oxidation. It is insoluble 
in water, but if the water contains air a small quantity of the 
arsenic is oxidized by the oxygen of the air to arsenious acid, which 
then dissolves in the water. 

When heated in the air arsenic volatilizes and burns, forming 
white fumes of arsenious anhydride. At the same time an unplea¬ 
sant garlic-like odour is noticed, which is perhaps peculiar to 
arsenic vapour, or is perhaps due to a lower oxide. Heated in 
oxygen, arsenic burns with a large evolution of light and heat, also 
producing arsenious anhydride. Finely powdered arsenic, when 
shaken into a jar of chlorine, catches fire and burns to arsenious 
chloride. Nitric acid oxidizes arsenic to arsenious or arsenic acid 
accordingly as the arsenic or nitric acid is in excess. In the same 
way concentrated sulphuric acid also oxidizes arsenic, and is itself 
reduced to a sulphurous anhydride. Hydrochloric acid scarcely acts 
upon it. 

Native arsenic is never pure. It may be obtained pure by 
sublimation from arsenical iron pyrites, which breaks up when 
heated into ferrous sulphide and metallic arsenic. 

Arsenic is chiefly used in the arts for hardening lead in the 
manufacture of shot. 


Arseniuretted Hydrogen ( Arsine ). 


231 


ARSENIURETTED HYDROGEN (Arsine). 

Composition : AsH 3 . 

This substance, corresponding to ammonia in composition, is 
the only known compound of arsenic and hydrogen. In conse¬ 
quence of the weak affinities of the two elements for one another 
it cannot be prepared directly from its constituents. 

Arseniuretted hydrogen is a colourless gas, with a repulsive 
odour resembling garlic, and only slightly soluble in water. It is 
condensed at-40 0 to a colourless liquid, and burns when ignited in 
the air with a livid flame forming water and arsenious anhydride. 
The specific gravity of the gas is 27, corresponding to the mole¬ 
cular formula : AsH 3 . Arseniuretted hydrogen is one of the most 
deadly poisons with which we are acquainted ; it is more poisonous 
than any other compound of arsenic. The chemist Gehlen lost his 
life by inhaling this gas, and fatal effects have also been produced 
in some other cases. Experiments with arseniuretted hydrogen 
must therefore be undertaken with great care, and the conditions 
under which the gas may be produced ought to be well known. 

Arseniuretted hydrogen is produced in all cases when nascent 
hydrogen and arsenic are brought together, especially in acid liquids. 
In considerable quantities it maybe obtained by acting on an alloy 
of zinc and arsenic (prepared by adding two parts of arsenic to three 
of melted zinc, and pouring the alloy into cold water) with dilute 
hydrochloric acid. The reaction is illustrated by the following 
equation :— 

As 2 Zn 3 + 3 S 0 2 |qh = 3S0 2 • 0 2 Zn + 2AsH 3 . 

Since commercial zinc nearly always contains small quantities 
of arsenic, the hydrogen which is prepared from zinc is nearly 
always contaminated with arseniuretted hydrogen. On the other 
hand, commercial hydrochloric and sulphuric acids also contain 
traces of arsenic as arsenious acid, and when such acids are used 
for the preparation of hydrogen, the nascent hydrogen reduces the 
arsenious acid to arseniuretted hydrogen :— 

As 2 O s + 6H 2 = 3 H 2 0 + 2 AsH 3 . 

On adding a few drops of a solution of arsenious acid to an appa¬ 
ratus evolving hydrogen from zinc and sulphuric acid, a mixture of 
hydrogen with sufficient arseniuretted hydrogen to exhibit its most 


232 Text-Book of Inorganic Chemistry. 

important properties is obtained. It is extremely dangerous to 
prepare the nearly pure gas from the alloy of zinc and arsenic. The 
production of the gas from a solution of arsenious acid and nascent 
hydrogen is used to detect extremely minute quantities of arsenic. 

The two constituents are so loosely combined in arseniuretted 
hydrogen, that the compound is decomposed at a low red-heat into 
arsenic and hydrogen. When led in a slow stream through a red- 
hot glass tube, the inner walls of the tube just beyond the heated 
portion become coated with a lustrous, nearly black mirror of re¬ 
duced arsenic. The mirror of arsenic may be readily observed 
when very minute quantities of the gas mixed with much hydrogen 
are passed through the hot tube. ’ The same decomposition takes 
place when a jet of arseniuretted hydrogen is inflamed in the air. 
By the high temperature of the flame the unburnt gas in its interior 
is decomposed into hydrogen and arsenic, and it is this finely 
divided arsenic which imparts to the flame its pale livid colour. 
That free arsenic is really present in the interior of the flame may 
be readily proved by pressing a piece of cold white porcelain, 
such as a porcelain dish, down upon the flame. The cold porcelain 
becomes coated with a round stain of black arsenic at the spot 
where the flame has touched it. The same phenomenon is produced, 
though less intense, if the gas is mixed with a large quantity of 
hydrogen. 

It is well known that wall-papers printed with colours (especially 
bright greens) containing arsenic exercise an injurious influence 
on persons living or sleeping in rooms so papered. The papers are 
fastened to the walls with starch paste, which easily ferments and 
sets free nascent hydrogen. And since the whole thickness of the 
paper is saturated with the paste, this nascent hydrogen comes 
directly into contact with the arsenic compounds,-and reduces them 
to arseniuretted hydrogen, which then poisons the air of the rooms. 
The quantity of the gas produced is of course very minute, but 
often sufficient to produce very injurious effects, and sometimes a 
faint unpleasant smell may also be observed. The sale of such 
papers ought therefore to be forbidden by law. 

That the arsenic and hydrogen are only feebly united in 
arseniuretted hydrogen is also shown by its general chemical 
behaviour. Chlorine decomposes the gas at once into arsenious 
chloride and hydrochloric acid, producing a large quantity of light 
and heat. This is easily shown by leading the gas into a cylinder 
containing chlorine. 


A rsenious Anhych'ide ( White A rsenic). 233 

Arseniuretted hydrogen led into a solution of silver nitrate 
produces a black precipitate of metallic silver, while both the 
hydrogen and arsenic are oxidized : — 

2 AsH 3 + i 2 N 0 2 - 0 Ag + 3 H 2 0 = 

I 2 N 0 2 - 0 H .+ i2Ag + As 2 0 3 . 

With copper sulphate it precipitates copper arsenide. 


COMPOUNDS OF ARSENIC WITH OXYGEN. 

Two of these compounds are known—arsenious anhydride and 
arsenic anhydride—both of which are soluble in water and then 
produce the corresponding acids. The former is a weak acid only 
slightly soluble in water, while the latter is easily soluble, and is a 
strong acid. The two oxides possess an analogous composition to 
that of phosphorous and phosphoric anhydrides, and arsenic acid, 
like common phosphoric acid, is tribasic. 

Just as when sulphur is burnt, sulphurous anhydride is produced 
and not sulphuric anhydride, so too the combustion of arsenic always 
produces arsenious anhydride and not arsenic anhydride. We can, 
however, easily oxidize arsenic, or the lower oxide, to arsenic acid 
by nitric acid, in the same manner as we obtain sulphuric acid from 
sulphur or its lower oxide. 


ARSENIOUS ANHYDRIDE. (White Arsenic.) 

Composition : As 2 0 3 . 

This oxide of arsenic is found in small quantities in nature as 
the mineral arsenic-bloom , and is produced on a large scale by 
roasting various minerals containing arsenic in a free supply of air. 
The volatile products so produced, which consist chiefly of sulphu¬ 
rous and arsenious anhydrides, are led through long passages and 
chambers, where the arsenious anhydride is deposited as a white 
crystalline powder. It is brought into trade partly in this form, 
and partly as a vitreous, amorphous solid. 

Arsenious anhydride is odourless, both in the solid and gaseous 
states. It possesses a faint sweetish taste. Two modifications of 




234 Text-Book of Inorganic Chemistry . 

this substance are known—the one amorphous and vitreous, the 
other crystalline. The former is prepared from the latter by 
subliming the white crystalline powder in upright iron retorts at 
as high a temperature as possible, when the vitreous form is de¬ 
posited in crusts on the neck of the retort. 

Vitreous arsenious anhydride is a transparent, lustrous, amor¬ 
phous solid, breaking with a conchoidal fracture. It has a specific 
gravity of 372, and gradually changes from the exterior to the 
interior into the crystalline modification and assumes the appear¬ 
ance of white porcelain. Pieces of vitreous arsenious anhydride 
which have been prepared for some time are found, when broken, 
to contain only a small nucleus of the transparent kind, all the 
rest having changed into the crystalline variety. 

The crystalline modification has a somewhat smaller specific 
gravity (3-62), and is less soluble in water than the vitreous kind. 
One part of the former requires 35 5 parts of water for its solution 
at the ordinary temperature, while one part of the latter dissolves 
in 108 parts of water; in boiling water more of each variety is 
dissolved. A hot saturated aqueous solution, on cooling, deposits 
arsenious anhydride in regular octahedra. The aqueous solution 
faintly reddens litmus paper. It is remarkable that, although 
arsenious anhydride is much less soluble in water than arsenic 
anhydride, it is still far more poisonous than this latter compound. 
Hydrochloric acid dissolves much larger quantities of arsenious 
anhydride than water. Nitric acid oxidizes it to arsenic acid, and 
is itself reduced to nitrous anhydride and nitric peroxide (p. 194). 

Arsenious anhydride sublimes when heated without melting, 
and is deposited on the cool walls of the vessel in brilliant trans¬ 
parent octahedra. It is dimorphous, and sometimes occurs in 
crystals belonging to the rhombic system. 

The vapour density of arsenious anhydride has been found 
to be 13-8, which corresponds to a molecule consisting of four 
atoms of arsenic united with six atoms of oxygen. It is, however, 
probable that if the vapour density were determined at a higher 
temperature it would be found to be only one-half as great, or, 
in other words, that the molecules of arsenious anhydride which 
exist at comparatively low temperatures would break up into 
molecules of half the weight, just as those of sulphur vapour, 
which contain six atoms at a temperature of about ioo° above its 
boiling point, break up at a higher temperature into molecules 
only one-third as heavy. 


A rsenious A nhydride ( White A rsenic). 235 

Arsemous Acid : As(OH) 3 , is probably contained in an aqueous 
solution of arsenious anhydride, but has not yet been prepared in the 
free state. It is a tribasic acid. With the alkalies it forms soluble 
salts ; its other compounds are insoluble or difficultly soluble in 
water, but are easily dissolved by acids, even by excess of arsenious 
acid itself. Its compounds are in general but little known ; many 
are unstable because of the weak affinities of the acid. The 
ammonium compound, obtained by saturating an aqueous solution 
of the acid with ammonia, loses all its ammonia on evaporation 
and finally leaves only arsenious anhydride behind. If to an 
aqueous solution of arsenious acid silver nitrate is added, no pre¬ 
cipitate of silver arsenite is formed ; because the nitric acid, which 
is set free on the production of this compound, keeps it in solution. 
But on the addition of one drop of ammonia, which neutralizes the 
nitric acid, a bright yellow precipitate is obtained, very similar to 
silver phosphate in appearance. It is also soluble in ammonia, 
and care must therefore be taken not to add too much of this 
reagent. Arsenious acid or its salts are at once distinguished from 
phosphoric acid or its salts by the production of a yellow precipi¬ 
tate of arsenious sulphide, when sulphuretted hydrogen is led 
through their slightly acid solution. 

The double compound of copper arsenite and copper acetate 
(Schweinfurt green), as well as an acid copper arsenite (Scheele’s 
green), are brilliant green pigments. Both are extremely poisonous, 
and are therefore but little employed; they should never be used 
for paper-hangings, textile fabrics, toys, and other similar articles. 

No substance has been for so long and so often the cause of 
death by poisoning both accidentally and intentionally as arsenious 
anhydride. Its external appearance, particularly its similarity to 
common flour, its slight taste, and lack of odour combine to render 
it particularly dangerous. Doses of 0-12 gramme (2 grains), or 
even less, often produce fatal effects. Chemists have, therefore, 
endeavoured to discover, firstly, some substance which, when taken 
internally, shall counteract the action of the poison— i.e. act as an 
antidote ; secondly, some method by means of which arsenic may 
be readily and certainly discovered in the bodies of persons who 
are supposed to have been poisoned by it. Both these objects have 
now been attained. 

Notwithstanding its poisonous properties, it is a remarkable 
fact that some animals—<?.£•. horses—can take large doses of arsenic 


236 Text-Book of Inorganic Chemistry. 

without any ill effects. On the contrary, it improves their general 
appearance and makes their coats more glossy. Doses, even 
as large as 15 grammes, have also been given to sheep without 
apparent injury. In the same way man himself may become 
accustomed to the use of arsenic. The so-called arsenic eaters 
of the Tyrol, who commence with small doses, at last take as much 
as \ gramme daily and even more. They thus become stouter, 
look healthier, and can ascend mountains more easily. Symptoms 
of arsenic poisoning only make their appearance when the use of 
the poison is discontinued. 

For more than 2,000 years it was vainly endeavoured to discover 
some substance which should act as an antidote to arsenic. In 
1834, however, Bunsen discovered such a substance in ferric 
hydrate, which, when taken soon enough, nearly always entirely 
counteracts the injurious effects of the poison. This discovery 
was not the result of accident, but was due to simple deductions 
from chemical facts. Bunsen found during his investigations on 
arsenious acid that ferric arsenite is quite insoluble in water; it 
was also known that substances insoluble in water and the gastric 
juice do not exercise any poisonous action on the animal body. 
From this he argued that if ferric hydrate in a suitable form were 
introduced into the stomach of persons who had taken arsenious 
acid, this insoluble and innocuous ferric arsenite would be formed 
and afterwards expelled from the body. 

His expectations were realised in a remarkable degree, and 
since then many valuable lives have been saved by the action of 
this antidote. It is only the freshly precipitated hydrate which 
possesses this action. If the hydrate-is kept for long, even when 
mixed with water, it undergoes a molecular change and becomes 
much weaker in its action. One of the best methods of administer¬ 
ing the antidote is as follows. A solution of ferric sulphate, pre¬ 
pared by oxidizing ferrous sulphate, is mixed with calcined mag¬ 
nesia (magnesium oxide). The substances then produced are the 
red-brown insoluble ferric hydrate and soluble magnesium sul¬ 
phate (Epsom salt), which remains in solution :— 

(S 0 ,) 3 0 6 Fe 2 + 3 MgO + 3 H 2 0 

= Fe 2 (OH) 6 + 3 S 0 2 . 0 2 Mg. 

The ferric hydrate combines with the arsenious acid, while the 
accompanying magnesium sulphate acts as a powerful purgative 
and rapidly removes the arsenic from the system. 


Detection of Arsenic in Cases of Suspected Poisoning. 237 


DETECTION OF ARSENIC IN CASES OF 
SUSPECTED POISONING. 

In cases of suspected poisoning by arsenic, it is first necessary 
to make a mechanical examination of the vomit or of the contents 
of the stomach to discover, if possible, white grains of unabsorbed 
arsenious anhydride. If such are found they are carefully collected 
and subjected to a special examination. 

One or two of these grains are placed at the end of a closed 
glass tube drawn out to a fine point, above which a small splinter 
of freshly glowed wood-charcoal is allowed to fall, as shown in 



Fig. 48. 

fig. 48. The splinter of charcoal is now gently heated in a gas 
flame until faintly glowing, and the tube is then slightly inclined 
so as to bring its extreme end into the flame. If the substance is 
arsenious anhydride it is converted into vapour, which on passing 
over the red-hot charcoal is reduced, its oxygen uniting with the 
charcoal, the reduced arsenic being deposited in the cooler parts 
of the tube above the charcoal as a lustrous black ring (fig. 49). 



Fig. 49. 

To be sure that this metallic mirror is really arsenic, the end of the 
tube is cut off when cold, the piece of charcoal allowed to drop 
out, and the mirror gently heated. The arsenic is then again con¬ 
verted into vapour, recombines with the oxygen contained in the 
warm air passing through the tube, and is deposited in the form 
of arsenious anhydride as a white crystalline coating in the upper 
and cooler parts of the tube. At the same time the garlic-like 
odour, produced when arsenic is volatilized, is noticed at the upper 
end of the tube. 

The white deposit of arsenious anhydride must dissolve in one 
drop of hydrochloric acid, and the solution, with the tube itself, 
when dropped into a test-tube containing sulphuretted hydrogen 
water, must give a bright yellow precipitate of arsenious sulphide. 




238 Text-Book of Inorganic Chemistry. 

Whether white grains of free arsenious anhydride have been 
found or not, it is usually necessary to look for the arsenic, which 
may be intimately mixed or chemically combined with the organic 
matter of the stomach and other organs. For this purpose it is 
first necessary to destroy the organic matter in the following 
manner. 

The organs are cut into small pieces, placed in a porcelain 
basin, and heated on the water-bath with chemically pure hydro¬ 
chloric acid, while small quantities of potassium chlorate are added 
from time to time. The chlorine which is liberated from the mix¬ 
ture of hydrochloric acid and potassium chlorate destroys the 
organic substances and oxidizes the arsenic to arsenic acid. The 
mixture must not be heated over the naked flame, or at least not 
boiled, for if this were done arsenic chloride would pass off with 
acid vapours and so be lost. As soon as the potassium chlorate 
has been completely destroyed, which may be easily recognized by 
the liquid no longer smelling of chlorine—the hydrochloric acid 
being of course in excess—the liquid is filtered and the residue 
well washed with hot water. All the arsenic, in the form of 
arsenic acid, is now contained in the solution, together with small 
quantities of organic compounds. From the clear liquid the 
arsenic is thrown down as arsenious sulphide by a long-continued 
stream of well-washed sulphuretted hydrogen. Arsenic acid, un¬ 
like arsenious acid, is not precipitated at once by sulphu¬ 
retted hydrogen, but only after some time. The first action of 
the gas is to reduce the arsenic acid to arsenious acid, free 
sulphur being at the same time precipitated. The yellow precipi¬ 
tate consists therefore of a mixture of sulphur and arsenious 
sulphide. 

As soon as the liquid smells strongly of sulphuretted hydrogen 
and has been allowed to stand for about twelve hours in a warm 
place, the precipitate—usually dark-coloured from organic im¬ 
purities—is collected on a small filter and well washed. The 
moist filter paper with the precipitate is then spread out inside a 
small porcelain dish, carefully dried, and moistened with pure con¬ 
centrated sulphuric acid. On gently warming the dish, the paper 
is completely charred and destroyed, as well as the traces of organic 
substances mixed with the precipitate ; a few drops of pure nitric 
acid or a crystal of nitre then completely oxidizes the arsenious 
sulphide again to arsenic acid. The filtered solution now con¬ 
tains the whole of the arsenic originally present in combination 


Detection of Arsenic in Cases of Suspected Poisoning . 239 

with oxygen as arsenic acid, and it now remains to prove in a 
certain manner the presence or absence of this substance. 

This is done by means of Marsh's apparatus , in which arsenic 
acid or arsenious acid is reduced to arseniuretted hydrogen, a gas 
easily yielding free arsenic when decomposed by heating. 

Marsh’s apparatus, so called after its discoverer, consists of a 
two-necked Woulffe’s bottle (fig. 50) furnished with a funnel-tube 
and a delivery-tube, both fitted gas-tight through good, sound 
corks. The short, bent delivery-tube is attached by a sound cork 
to a piece of difficultly fusible gas (called combustion-tube) which 
has been previously narrowed in several places, and of which the 
end is bent upwards at right angles, and terminates in a small jet. 
The front part of this tube is packed loosely with cotton-wool or 



Fig. 50. 


glass-wool, to retain any liquid which might be mechanically 
carried over from the Woulffe’s bottle by the gas. The tube is 
supported in several places so that it may not bend when after¬ 
wards heated. 

After a quantity of granulated zinc, free from arsenic, has been 
introduced into the Woulffe’s bottle and the apparatus arranged 
as described, a cold mixture of pure sulphuric acid with about 
eight times its volume of water, prepared beforehand, is poured 
into the flask through the funnel tube. The hydrogen, which is at 
first evolved slowly, but afterwards more rapidly, is allowed to 
pass through the apparatus until all the air has been expelled. 
The gas is then ignited at the jet, and the tube strongly heated in 
several places by separate powerful burners. 
























240 Text-Book of Inorganic Chemistry. 

This preliminary experiment is to decide whether the substances 
employed—the zinc and the sulphuric acid—are absolutely free from 
arsenic. If either contains even traces of arsenic, this substance 
is always converted by the nascent hydrogen into arseniuretted 
hydrogen. This latter gas, even when mixed with a large excess 
of hydrogen, is decomposed when heated to redness in the tube 
into arsenic and hydrogen, and the former is deposited as a brilliant 
black ring or mirror on the cooler parts of the tube just beyond 
the portions heated. 

Should the stream of gas be rapid, small quantities of arseniu¬ 
retted hydrogen may remain undecomposed, but are at once de¬ 
tected by pressing a cold white porcelain dish on the jet of burning 
gas. If arsenic is present, that part of the porcelain dish in 
contact with the flame becomes covered with a black stain of 
arsenic. 

If after ten minutes of continued heating no trace of a mirror 
is formed in the tube, and the porcelain dish when pressed down in 
the flame remains perfectly white, it is absolutely certain that both 
the zinc and the sulphuric acid contain no arsenic. 

The dilute sulphuric acid solution which is to be tested for 
arsenic is now gradually poured into the Woulffe’s bottle, while the 
tube is continuously heated in several places and the jet at the end 
kept ignited. The rate at which the gas is evolved is to be regu¬ 
lated by pouring in acid or by placing the bottle in cold water : it 
was for this purpose previously placed in an empty basin. The 
rapidity of the stream of gas is known by the height of the flame at 
the jet; it should be regulated until the flame is about three centi¬ 
metres (one inch) high. If arsenic is present, as many of the 
mirrors as possible, both in the tube and on porcelain dishes, 
should be prepared in order to prove by further experiments that 
they really consist of arsenic. And when cold the glass tube is 
afterwards cut up into as many pieces as it contains mirrors. 

Even with all the precautions which have just been enumerated, 
there is still a possibility of error. The oxides of antimony, like 
those of arsenic, are also reduced by nascent hydrogen, producing 
gaseous antimoniuretted hydrogen, which, like the arsenic com¬ 
pound, is decomposed at a red heat into its constituents—antimony 
and hydrogen. If, therefore, the suspected liquid contains no 
arsenic but antimony, mirrors of antimony are obtained in the 
Marsh’s apparatus which can scarcely be distinguished by the naked 
eye from those of arsenic. 


Detection of Arsenic in Cases of Suspected Poisoning. 241 

And when it is remembered that in cases of suspected poison¬ 
ing an emetic is nearly always given to remove the poison as quickly 
as possible from the stomach, and that one of the most powerfully 
acting of these substances is tartar emetic (a compound of tartaric 
acid, potassium, and antimony), it will be at once seen that on 
treating the contents of the stomach by the above process, mirrors 
of antimony may be obtained indistinguishable from those of 
arsenic without further experiments. 

To be quite sure that the mirrors (if obtained) are arsenic and 
not antimony, the following experiments may be tried :— 

1. One of the pieces of glass tube containing a mirror is im¬ 
mersed in a freshly prepared solution of sodium hypochlorite. If 
the mirror is of arsenic it disappears almost immediately, while if of 
antimony it remains unchanged in the liquid for more than twenty- 
four hours. 

2. A second piece of tubing containing a mirror is inclined at 
an angle and gently heated. Both antimony and arsenic are thus 
oxidized, the former to antimony tetroxide, the latter to arsenious 
anhydride, which are deposited on the upper and cooler parts of 
the tube. Both these white deposits dissolve readily in a drop of 
warm hydrochloric acid, but on dipping them into sulphuretted 
hydrogen water, the antimony gives an orange-coloured precipi¬ 
tate of antimonous sulphide, while the corresponding sulphide of 
arsenic obtained in the same way is bright yellow. 

3. The mirrors obtained by depressing a piece of cold porcelain 
on a jet of the burning gas are treated with yellow ammonium sul¬ 
phide (containing dissolved sulphur). Both arsenic>and antimony 
dissolve (the latter more quickly), and so produce compounds of am¬ 
monium sulphide with arsenic and antimony sulphides respectively. 
If, then, both are evaporated to dryness, yellow arsenic sulphide is 
produced from the former and orange antimony sulphide from the 
latter. The orange antimony sulphide is at once dissolved by a 
few drops of warm hydrochloric acid, but the yellow arsenic sul¬ 
phide remains unchanged. 

Such investigations to discover the presence or absence of 
arsenic are not of themselves difficult. But they require great 
care and experience on the part of the chemist; he must be abso¬ 
lutely certain (i.) that small quantities of arsenic have not escaped 
him, (ii.) that arsenic was not contained in tho reagents or vessels 
employed and so have been found when not really present in the 
original substance ; and, finally, that antimony has not been mis- 

R 


242 Text-Book of Inorganic Chemistry. 

taken for arsenic. In criminal cases, where his decision may be a 
question of life or death, the investigations should be conducted 
with even greater care, and only chemists of long experience should 
be employed. 

Among the minor precautions which must be remembered and 
attended to in such investigations are the following. 

Pure diluted sulphuric acid and pure zinc evolve pure hydrogen 
without a trace of sulphuretted hydrogen, even when the liquid be¬ 
comes heated. But if concentrated sulphuric acid is added by a 
funnel tube to a mixture of zinc and water, the hydrogen is found 
to be mixed with sulphuretted hydrogen, the latter gas being pro¬ 
duced by the action of the nascent hydrogen upon the sulphurous 
anhydride formed under these circumstances (p. 160). If arsenious 
acid is present in this liquid it is at once converted into arsenious 
sulphide, upon which nascent hydrogen has no action. 

If, therefore, concentrated acid is poured into the Marsh’s appa¬ 
ratus— e.g. to accelerate the evolution of gas—and the solution to 
be tested for arsenic then added, even if the latter substance were 
present, no mirror or only a minute one of arsenic would be pro¬ 
duced, because the arsenic would remain behind in the flask as 
arsenious sulphide. 

For the same reason granulated zinc which has lain in the 
laboratory for some time and of which the surface has become 
covered with a thin coating should first be digested with dilute 
sulphuric acid and washed. This coating is not always simply zinc 
oxide, but may contain zinc sulphide, which would produce sul¬ 
phuretted hydrogen in the Marsh’s apparatus and lead to the forma¬ 
tion of arsenious sulphide. 

Since sulphurous acid, in acid solutions, is reduced by nascent 
hydrogen to sulphuretted hydrogen, care must be taken that the 
liquid before introduction into the Marsh’s apparatus is perfectly 
free from this substance. A mere trace of sulphurous acid would 
form arsenious sulphide, and so be a source of error. 

If the liquid to be tested contains organic matter suspended or 
dissolved when introduced into the Marsh’s apparatus, it might 
be volatilized or particles of it might be carried over mechanically 
by the stream of gas, and become carbonized at the heated portions 
of the tube. A dark lustrous deposit of carbon might thus be 
produced similar to the arsenic mirror. It is, for this reason, 
absolutely necessary to remove the whole of the organic matter 
by treating the original substance with hydrochloric acid and 


Detection of Arsenic in Cases of Suspected Poisoning. 243 

potassium chlorate, and by acting on the sulphuretted hydrogen 
precipitate with concentrated sulphuric acid. 

In some cases the dried substance might be at once charred 
with concentrated sulphuric acid to destroy organic matter, without 
first acting upon it with hydrochloric acid and potassium chlorate. 
This, however, is only permissible when the substance is free from 
common salt or other metallic chlorides. In such cases hydro¬ 
chloric acid is formed, which unites with the arsenious anhydride 
to form volatile arsenious chloride, and this passes away with the 
vapours, and is lost. And as the organs of the human body nearly 
all contain common salt, this charring with sulphuric acid is in¬ 
admissible. 

It is of course of the utmost importance to test the various 
substances employed in the investigation for arsenic. The sul¬ 
phuric acid, zinc, hydrochloric acid, potassium chlorate, the nitric 
acid or nitre and even the filter papers and water must all be 
previously brought into the Marsh’s apparatus for this purpose. 
It is also better in all important (criminal) cases to employ exclu¬ 
sively new vessels and apparatus which have never previously been 
used for chemical purposes. 

A further complication arises when the substance to be tested 
for arsenic contains antimony. Mirrors are then produced in the 
Marsh’s apparatus, but no definite result can be obtained, espe¬ 
cially when only relatively small quantities of arsenic are present. 
If the mirror does not apparently dissolve when dipped into a 
solution of sodium hypochlorite, it does not follow that arsenic is 
absent. 

We owe to Fresenius a very simple and exact method of de¬ 
tecting arsenic in presence of much larger quantities of antimony. 
This method depends upon the fact that arsenic is easily volatilized 
at a low red heat, but antimony only at a much higher temperature, 
and that the sulphides of both metals are easily reduced when 
heated with potassium cyanide, which is then converted into potas¬ 
sium sulphocyanate. The practical details are as follows. 

The substance to be tested for arsenic is treated as described 
above, except that finally the antimony and arsenic are precipitated 
as sulphides by a continued stream of sulphuretted hydrogen, and 
the sulphides washed and dried. The dried sulphides are then 
rapidly mixed in a warm mortar with about four times as much of 
a mixture of 1 part of potassium cyanide and 3 parts of dry sodium 
carbonate. This mixture is introduced quickly (to avoid abstraction 



244 Text-Book of Inorganic Chemistry . 

of moisture from the air) into a short piece of glass tubing open at 
both ends and this slipped into a piece of combustion tube drawn 
out at one end. A carbonic acid apparatus is connected with this 
tube as shown in fig. 51. 

As soon as all the air is expelled the stream of carbonic acid is so 
regulated that it only passes quite slowly. Various positions of the 
narrow portion are now heated to low redness, and then the mixture 
of the two sulphides, with sodium carbonate and potassium cyanide. 



tively. The latter remains with the mixture, while the former is 
volatilized and deposited beyond the heated portions of the narrow 
tube as ordinary arsenic mirrors which exhibit all the properties of 
this element. At the same time a small quantity of arsenic is 
carried over with the carbonic acid, imparting to the gas issuing 
from the end of the tube the well-known intense garlic odour of 
arsenic. 


ARSENIC ANHYDRIDE : As 2 0 5 , and ARSENIC 
ACID: AsO•(OH) 3 . 

On the combustion of arsenic in oxygen, the lower oxide, arse- 
nious anhydride, and not the higher, arsenic anhydride, is always 
produced, because the latter compound is decomposed at the tern- 





















Arsenic Acid. 


245 


perature of combustion of arsenic into arsenious anhydride and 
oxygen. We have already seen (p. 156) that for the same reason 
sulphur always burns to sulphurous anhydride and not to sulphuric 
anhydride. The oxidation can, however, be readily effected in 
what is called the wet way by nitric acid, which easily gives up 
oxygen to many substances. 

When arsenious anhydride is heated with nitric acid large 
quantities of red fumes consisting of nitric peroxide and nitrous 
anhydride are evolved ; and if the strongly acid liquid is evapo¬ 
rated down to a syrup in order to expel the excess of nitric acid, 
crystals of arsenic acid united with water: 2AsO(OH) 3 + H 2 0 , 
separate out on cooling in rhombic plates. These crystals are 
very deliquescent; they melt at ioo° and lose their water of crystal¬ 
lization, leaving pure arsenic acid. 

Arsenic acid so obtained is a strong acid, and, like ordinary 
phosphoric acid, is tribasic. It also loses water when heated to 

^ (AsO(OH) 2 a , 

about 18o°,forming ietrabasic /yroarsemcacid : <J |AsO(OH)„ 

a higher temperature, about 200°, a further quantity of water is 
given off and the acid becomes monobasic metarsenic acid : 
AsCVOH, which, finally, at a low red heat is converted into arsenic 
anhydride : As 2 O s , thus :— 

2 As 0 2 - 0 H - h 2 o - As 2 0 5 = O 

The anhydride is a white amorphous substance, only difficultly 
soluble in water, but which is gradually converted into ordinary 
tribasic arsenic acid when allowed to stand for some time in contact 
with water. At temperatures above a red heat arsenic anhydride 
breaks up into arsenious anhydride and oxygen. 

Even at ordinary temperatures arsenic acid in solution readily 
gives up oxygen to many substances, as, for example, sulphuretted 
hydrogen, sulphurous anhydride, stannous chloride, &c., and is 
reduced to arsenious acid. Nascent hydrogen evolved in acid 
liquids reduces arsenic acid, as well as arsenious acid, to arseniu- 
retted hydrogen. In consequence of the readiness with which 
arsenic acid parts with oxygen, it is largely employed in the arts— 
principally to oxidize aniline, and to convert it into the brilliant 
aniline colour known as magenta or fuchsine. 

Common or tribasic arsenic acid yields, like common phos¬ 
phoric acid, three series of salts—the normal, monacid, and 
diacid salts. The two former are mostly insoluble, or difficultly 


246 Text-Book of Inorganic Chemistry. 

soluble, in water. Normal silver arsenate : AsO(OAg) 3 , is thrown 
down as a chocolate-coloured precipitate on mixing solutions of 
the sodium salt and silver nitrate. It is easily soluble in nitric 
acid, and is at once distinguished from yellow silver arsenite 
by its colour. The monacid sodium arsenate corresponding to 
common sodium phosphate has the composition : — 

As ° { oh* 1 ^ + 12 H 2°> 

and, like the phosphate, is obtained in fine clear crystals when its 
solution is slowly evaporated. 

Arsenic acid, like phosphoric acid, also produces a yellow 
crystalline precipitate with a nitric acid solution of ammonium 
molybdate. And the compound of arsenic acid corresponding to 
ammonium-magnesium phosphate (p. 218), of the composition : 

AsO | is, like the phosphate, insoluble in water containing 

ammonia, and is therefore used for the quantitative determination 
of arsenic. 

Arsenic acid, like arsenious acid, is precipitated as a yellow 
sulphide by sulphuretted hydrogen. This precipitation is not, 
however, produced at once, as is the case with arsenious acid, but 
only after some time, and then consists not of arsenic pentasulphide 
but of a mixture of the trisulphide and sulphur. By warming the 
solution, which assists the reduction of the arsenic acid to arse¬ 
nious acid, the precipitation may be accelerated. Arsenic acid is 
poisonous, but, notwithstanding its much greater solubility in 
water, is far less so than arsenious acid. 


COMPOUNDS OF ARSENIC WITH THE HALOGENS. 

Arsenic, like phosphorus, combines directly with the halogens, 
but of these compounds only those are known which contain triad 
arsenic. Compounds containing one atom of arsenic united with 
five atoms of chlorine, bromine, or iodine have not yet been pre¬ 
pared. 



Compounds of Arsenic with the Halogens. 247 

Arsenious Chloride : AsC 1 3 , is a colourless, volatile, oily 
liquid, boiling at 130°, and with a specific gravity of 2’2. It is 
easily decomposed by water into hydrochloric and arsenious acids, 
and is therefore very poisonous. Its vapour density is 6*3, corre¬ 
sponding to the molecular weight represented by the formula : 
AsC 1 3 . 

Arsenious chloride is produced by acting upon arsenic with 
dry chlorine, or better, by heating an intimate and dry mixture of 
4 parts arsenious anhydride and 7 parts common salt with an 
excess of concentrated sulphuric acid in a retort provided with 
a receiver :— 

As 2 0 3 + 6NaCl + 6 S 0 2 ( 0 H) 2 = 

2AsC1 3 + + 3H 2 0. 

The arsenious chloride then distils over, and acid sodium sulphate 
remains behind. The water, which is produced at the same time, 
is retained by the excess of sulphuric acid. 

A simpler and better method of preparation consists in heating 
a mixture of powdered arsenic (1 part) with dry mercuric chloride 
(io| parts) in a retort. The mercuric chloride then parts with one 
half of its chlorine, and is converted into mercurous chloride : 

2 As + 6HgCl 2 = 2 AsC 1 3 + 3Hg 2 Cl 2 . 

The chloride is purified by redistillation ; it often possesses a 
violet colour. 


Arsenious Bromide : AsBr 3 , is also produced by the direct 
union of its elements, best by dissolving the dry bromine m per¬ 
fectly dry carbon disulphide and then adding small quantities of 
powdered arsenic. By slow evaporation of the clear liquid poured 
off from the excess of arsenic in a current of dry air, arsenious 
bromide crystallizes out. It deliquesces when exposed to the air, 
and is easily decomposed by water, like the chloride, into hydro- 
bromic and arsenious acids. 


Arsenious Iodide : Asl„ is prepared m a similar manner to 
that employed for the bromide, and crystallizes on evaporation of 
the solvent as brilliant red tablets. This compound is soluble m 
alcohol, and maybe crystallized from its solution in this liquid. 
It is dissolved unchanged by cold water, but. if the solution is 


248 Text-Book of Inorganic Chemistry. 

warmed decomposition ensues with formation of hydriodic and 
arsenious acids. Arsenious iodide is employed in medicine as a 
remedy for cancer. 

Arsenious Fluoride : AsF 3 , is produced as a colourless liquid 
by distilling equal parts of finely powdered fluor-spar and arsenious 
anhydride, with five parts concentrated sulphuric acid. It boils 
at 6o° and fumes strongly in the air. Water first dissolves it, 
but decomposition into hydrofluoric and arsenious acids soon 
occurs. 


COMPOUNDS OF ARSENIC AND SULPHUR. 

We are acquainted with three compounds of arsenic and 
sulphur—viz., a disulphide : As 2 S 2 ; a trisulphide : As 2 S 3 ; and a 
pentasulphide : As 2 S 5 ; of which the two former occur in nature. 

Arsenic Bisulphide : As 2 S 2 . This compound is found in nature 
as fine red rhombic prisms as the mineral realgar. It may also be 
prepared from its constituents by heating a mixture of arsenic and 
sulphur in the proper proportions, or by heating a mixture of 
arsenious anhydride and sufficient sulphur, sulphurous anhydride 
being then evolved. It is insoluble in water, and is first decom¬ 
posed and then dissolved by alkaline sulphides. 

The substance which occurs in trade as a red vitreous mass 
under the name of realgar , ruby-sulphur, red arsenic glass , and 
which is used to some extent as a pigment, is not pure arsenic 
disulphide. It is prepared by distilling a mixture of arsenical 
pyrites (FeAsS) with common pyrites (FeS 2 ), and contains arsenic 
trisulphide, arsenious anhydride, or even free arsenic. This crude 
disulphide is employed for the manufacture of the so-called Bengal 
fire. A mixture of 24 parts nitre, 7 parts sulphur, and 2 parts ruby- 
sulphur burns with a penetrating white light when ignited. 

The corresponding oxide, of the composition : AsO or As 2 0 . has 
not yet been prepared. 2 ’ 


Arsenic Trisulphide (Arsenious Sulphide) : As 2 S 3 , occurs 
in nature as yellow rhombic prisms and is then called orpi- 
ment. Artificially it is prepared by heating together arsenic and 
sulphur in the proper proportions, or by passing a stream of sul- 


Compounds of Arsenic and Sulphur. 249 

phuretted hydrogen though an aqueous solution of arsenious 
acid, in the presence of hydrochloric acid. In the latter case the 
trisulphide is obtained on drying as a pale yellow, amorphous 
powder. 

Arsenic trisulphide melts when heated to form a red liquid, 
which solidifies to a red semi-transparent glass. It is more trans¬ 
parent and gives a lighter coloured powder than the disulphide. 
Water and hydrochloric acid do not dissolve it, nitric acid oxidizes 
it to arsenic acid and sulphuric acid; it is also converted into 
arsenic acid by concentrated sulphuric acid, sulphurous anhydride 
being then set free. When mixed with potassium cyanide and 
heated in a glass tube it is reduced to arsenic, which sublimes, while 
the sulphur unites with the potassium cyanide forming potassium 
sulphocyanate, thus :— 

3KCy + AsjS 3 = 2As + 3KCyS. 

Arsenic trisulphide unites with basic sulphides to form sulpho- 
salts. The majority cf these compounds are insoluble in water, 
but those containing the alkaline sulphides are soluble. It is at 
once dissolved by a cold solution of an alkaline sulphide, from 
which solution hydrochloric acid again precipitates the whole of 
the arsenic trisulphide. The potassium salt has the composition, 
AsS-SK, and its decomposition by dilute acids is shown in the 
following equation :— 

2AsS-SK + 2HCI = 2KCI + As 2 S 3 + H 2 S. 

On mixing a solution of potassium sulpharsenite with solutions 
of silver, copper, or lead nitrate, a precipitate, usually coloured, of 
the corresponding silver, copper, or lead compound is produced, 
thus :— 

AsS-SK + N 0 2 • O Ag = N 0 2 - 0 K + AsS-SAg. 

Caustic alkalies and even the alkaline carbonates also dissolve 
arsenic trisulphide, producing a mixture of an arsenite with a sulph¬ 
arsenite. With caustic potash, for example, the reaction is as 
follows :— 


2As 2 S 3 + 4KOH 


AsO-OK + 3 AsS-SK + 2 H 2 0 , 


250 


Text-Book of Inorganic Chemistry , 

Arsenic Pentasulphide (Arsenic Sulphide) : As. 2 S 5 , is a yel¬ 
low powder closely resembling the trisulphide in external properties, 
but which does not occur in nature. It might be supposed that 
sulphuretted hydrogen when led into a solution of arsenic acid 
would give a precipitate of arsenic pentasulphide, but this is not 
confirmed by experiment. The gas may be led into the acid liquid 
for a considerable time without producing any precipitate or even 
turbidity. Only after some time a yellow precipitate is gradually 
produced, which is not the pentasulphide, but a mixture of the tri¬ 
sulphide with free sulphur. That this is really the case may be 
proved by digesting the dried precipitate with carbon disulphide, 
which dissolves out the free sulphur, while yellow arsenic trisul¬ 
phide remains behind. 

Arsenic pentasulphide is best prepared by saturating a solution 
of potassium arsenate with sulphuretted hydrogen, and then decom¬ 
posing the potassium sulpharsenate, so obtained, with dilute hydro¬ 
chloric acid. The following equations represent the process:— 

(i.) AsO(OK) 3 + 4H0S = AsS(SK) 3 + 4 H 2 0 . 

Potassium Potassium 

ar.-enate sulpharsenate 

(ii.) 2 AsS(SK) 3 + 6 HC 1 = As,S 5 + 6 KC 1 + 3H 2 S. 

Arsenic pentasulphide combines with other sulphides and forms 
the sulpharsenates, corresponding to the above potassium com¬ 
pound. 


ANTIMONY. 

Atomic Weight : 120 .—Chemical Symbol: Sb. 

Antimony very seldom occurs free in nature, but is usually found 
combined with sulphur. Its commonest form is the trisulphide 
(Sb 2 S 3 )—the mineral called grey antimony ore or antimonite. This 
trisulphide also occurs in nature in combination with other sulphides 
as a sulpho-acid of which the minerals chalcostibite (Cu 2 S,Sb 2 S 3 ) 
and dark red silver ore or pyrargyrite (3Ag 2 S,Sb 2 S 3 ) are the com¬ 
monest. In combination with nickel it is found as breithauptite 
(NiSb), and with silver as dyscrasite (Ag 4 Sb). 

In its physical properties antimony so closely resembles the 
metals, especially bismuth, that it is sometimes included in this 
group of elements. Its chemical properties and the compounds 



A ntimony. 


251 


which it forms show, however, that it is much more closely allied 
to the nitrogen group of the non-metals. It is a lustrous crystal¬ 
line solid of a bluish-white colour, with a specific gravity of 67 ; 
melts at about 440°, and crystallizes on cooling in rhombohedra. 
When slowly cooled its fracture shows large crystalline laminae, but 
when quickly cooled the fracture is granular. It is volatilized at a 
bright red heat, and may be distilled at a white heat in a stream of 
hydrogen gas. 

The antimony of commerce is obtained almost exclusively from 
the trisulphide—grey antimony ore. The ore melts at a low tem¬ 
perature and can thus be easily separated from earthy impurities 
accompanying it. The purified ore which forms a dark grey, 
lustrous, and crystalline mass, is then heated to redness in a cru¬ 
cible with 42 per cent, of wrought iron scrap. By this means 
ferrous sulphide and a regulus of antimony are obtained, and the 
latter separates better from the slag if some dried sodium sulphate 
(10 parts) and charcoal powder (3 parts) are added before heating. 
These two substances form carbonic oxide and sodium sulphide, 
the latter then uniting with the ferrous sulphide to form an easily 
fusible slag. 

Another method of extracting antimony consists in roasting the 
ore in a reverbatory furnace and then reducing the antimony oxide 
so formed with charcoal and sodium carbonate. 

Antimony so prepared is never pure ; it usually contains lead, 
iron, and copper, and nearly always traces of arsenic, from which it 
must be completely purified before it can be employed for pharma¬ 
ceutical preparations, such as tartar emetic, antimony chloride and 
oxide. 

Crude antimony is purified by fusing 16 parts of it with 2 parts 
of dry sodium carbonate and 1 part of antimony trisulphide in 
a Hessian crucible for one hour. The regulus so obtained is 
then again fused for the same length of time with parts of 
sodium carbonate, and finally a third time with 1 part of sodium 
carbonate and a little nitre. The arsenic which was not converted 
in the first fusion into sodium sulpharsenite is oxidized to potas¬ 
sium arsenate by the nitre, leaving the antimony then free from 
arsenic. 

Antimony remains unaltered when exposed to dry air. When 
small quantities are heated on charcoal before the blowpipe to 
above its boiling-point antimony burns, forming white fumes of 
antimonous oxide which are partly deposited on the charcoal. If the 


252 Text-Book of Inorganic Chemistry. 

fused metal is then allowed to cool it becomes covered with a net¬ 
work of transparent crystals of the oxide. 

Powdered antimony burns brilliantly in chlorine gas, forming 
one of its chlorides. It also unites directly with sulphur. 

Hydrochloric or dilute sulphuric acid have no action on anti¬ 
mony ; nitric acid easily oxidizes it to one of the oxides of antimony, 
forming a white powder insoluble both in water and nitric acid. 
Aqua regia alone dissolves it, producing either the tri- or penta- 
chloride. Both these chlorides dissolve in hydrochloric acid, but 
are precipitated by water as oxychlorides. Tartaric acid prevents 
the precipitation of these oxychlorides. 

Antimony is used for many other purposes besides the prepara¬ 
tion of useful medicines. It enters into the composition of some 
important alloys, such as type-metal (antimony, lead, and tin) and 
Britannia-metal (antimony, tin, and zinc). 


ANTI MON IURETTED HYDROGEN (stibine). 

Composition : SbH 3 . 

This gaseous compound very closely resembles the correspond¬ 
ing arseniuretted hydrogen, but is less poisonous. Antimoniuretted 
hydrogen is prepared, in the same way as the arsenic compound, 
from an alloy of antimony and zinc, or better, by acting upon a 
compound of antimony and potassium with dilute sulphuric or 
hydrochloric acid. It is, further, always produced when nascent 
hydrogen comes into contact with a soluble antimony compound in 
an acid solution. 

Antimoniuretted hydrogen is a colourless gas without odour, is 
decomposed at a low red heat, and burns in the air with a greenish 
flame, forming antimonous oxide and water. A piece of cold white 
porcelain depressed on this flame receives a black stain of free 
antimony. We have already described how this antimony stain 
may be distinguished from a corresponding one of arsenic (p. 241). 
It may be mentioned that when antimoniuretted hydrogen is passed 
into a solution of silver nitrate a black precipitate of a compound 
of silver and antimony is formed, while arseniuretted hydrogen 
under the same conditions precipitates metallic silver and the 
arsenic becomes oxidized to arsenious acid. 



Compounds of Antimony and Oxygen. 


253 


COMPOUNDS OF ANTIMONY AND OXYGEN. 

Two compounds of antimony and oxygen are known, resem¬ 
bling in composition the two oxides of arsenic—viz. antimony tri¬ 
oxide, or antimonous anhydride (Sb 2 0 3 ), corresponding to arsenious 
anhydride (As 2 0 3 ), and antimonic, anhydride (Sb 2 0 5 ), corresponding 
to arsenic anhydride (As 2 O s ). Besides these, a third compound 
also exists, having the composition : Sb 2 0 4 , which may, however, 
be considered as a compound of the other two, thus : Sb 2 0 3 , 
Sb 2 0 5 or (Sb 0 2 ) 0 (Sb 0 ). 

The general chemical affinities of antimony are considerably 
weaker than those of arsenic, and antimony trioxide is in fact more 
of a base than an acid, but is so weak a base that its compounds 
with acids are immediately decomposed by water. 


ANTIMONY TRIOXIDE. 

Compositio 7 i : Sb 2 0 3 or (SbO) 2 0 . 

This compound occurs in nature in two distinct crystalline 
forms. Firstly, as rhombic prisms in the mineral valentinite , and 
secondly, in regular octahedra as senarmontile. Both forms may 
be obtained by burning antimony in the air and condensing the 
white vapours on a cold body. Both are isomorphous with the two 
forms in which arsenious anhydride crystallizes, and the two sub¬ 
stances are therefore said to be isodimorphous. 

In the wet way antimony trioxide may be obtained by acting 
upon finely powdered antimony with dilute nitric acid, or by pre¬ 
cipitating a solution of antimony trichloride with water and washing 
the white oxychloride with sodium carbonate, when white insoluble 
antimony trioxide remains behind. 

The dried oxide becomes yellow when heated but is again 
colourless when cold; when more strongly heated it melts, and at 
still higher temperatures it sublimes, taking up oxygen from the 
air and becoming converted into antimony tetroxide. 

Antimony trioxide is insoluble in nitric acid and in dilute 
sulphuric acid, but is dissolved by hydrochloric acid, concentrated 



254 Text-Book of Inorganic Chemistry. 

sulphuric acid, 1 or tartaric acid, as well as by acid potassium tar¬ 
trate. In the last case a soluble crystalline compound is formed, 
called tartar emetic —potassium antimony tartrate. 

With strong bases antimony trioxide behaves as an acid. It 
dissolves, for example, in strong caustic soda, and the solution on 
cooling deposits crystals of sodium antimonite, of the composition : 
SbO-ONa + 3H 2 0. In most of its compounds the trioxide takes 
the form of the monad radical : SbO. 

The compound obtained on precipitating antimony trichloride 
with sodium carbonate is not antimony carbonate but simply the 
hydrated trioxide, which loses its water on boiling. Antimony tri¬ 
oxide is too weak a base to combine with carbonic acid. 


ANTIMONIC ANHYDRIDE. (Antimony 

Pentoxide.) 

Composition : Sb 2 0 5 or (Sb 0 2 ) 2 0 . 

Antimonic anhydride, obtained by gently heating either of the 
antimonic acids, is a bright yellow powder insoluble in water, and 
which when strongly heated does not melt but decomposes into 
antimony tetroxide and free oxygen. 

Corresponding to this compound are two antimonic acids, one 
monobasic and one tetrabasic. 

Antimonic Acid : Sb 0 2 * 0 H + H 2 0 , is a white powder, which, 
though scarcely soluble in water, reddens litmus paper. It may be 
obtained by acting on powdered antimony with aqua regia contain¬ 
ing an excess of nitric acid, or by heating the metal for a long time 
with strong nitric acid. On heating a mixture of powdered anti¬ 
mony with four times its weight of nitre a deflagration takes place 
and the saline mass, when afterwards extracted with luke-warm 
water, leaves a white powder of potassium antimonate : Sb 0 2 -OK, 
This compound is only slightly soluble in cold water and can only 
be slowly dissolved by continued boiling. If the solution so 
obtained is then evaporated, the salt remains behind as a gummy 

1 From the solution of antimony trioxide in concentrated sulphuric acid the 
salt, antimonous sulphate : (S0 2 ) 3 O e Sb 2 , in which the oxide is a base, crystal¬ 
lizes out. This salt is decomposed by water into free sulphuric acid and basic 
sulphates.—E d. 



Antimonic Acid. 


255 

mass. Nitric acid decomposes it, giving a white precipitate of 
antimonic acid. 

This form of antimonic acid is soluble in concentrated hydro¬ 
chloric acid, and easily soluble in caustic potash, but it is not dis¬ 
solved by ammonia. Its salts, even those of the alkalies, are 
mostly insoluble or -difficultly soluble in water and are easily de¬ 
composed even by weak acids. 


Metantimonic Acid : 1 H 4 Sb 2 0 7 = O j 

{ SbO(OH) 2 

This tetrabasic acid resembles the preceding monobasic acid 
in its external properties ; it is, however, more soluble in water and 
acids and in ammonia. It is formed when antimony pentachloride 
is decomposed by water :— 

2SbCl 5 + 7H 2 0 = H 4 Sb 2 0 7 + 10HCI, 

or by acting on a metantimonate with hydrochloric acid. In both 
cases it separates as an amorphous precipitate. In combina¬ 
tion with potassium it may be easily obtained by fusing ammo¬ 
nium or potassium antimonate with three times its weight of 
caustic potash. The saline mass is then dissolved in water and 
evaporated down, when potassium metantimonate separates out 
as deliquescent crystals on cooling. A small quantity of cold 
water decomposes this—the normal salt—into free alkali and the 
diacid compound (H 2 K 2 Sb 2 0 7 + 6 H 2 0 ) which remains behind as 
a granular mass. Hydrochloric acid separates metantimonic acid 
from both salts. 

Metantimonic acid possesses the remarkable property of form¬ 
ing a compound with soda insoluble in water. If a solution of 
sodium chloride or of any other sodium salt is added to the above- 
named solution of potassium metantimonate, containing free alkali, 
a white precipitate of sodium metantimonate is formed. On 
account of this reaction potassium metantimonate is sometimes 
used as a reagent for sodium compounds. It is not, however, much 
employed for this purpose, as the solution of the potassium salt 
gradually changes into the ordinary antimonate, which does not 
produce an insoluble salt with sodium compounds. 

1 By analogy with the acids of phosphorus, this acid should be called 
pyrantimonic acid and the preceding compound metantimonic acid. The 
names given above are, however, those in general use among chemists.— Ed. 


256 Text-Book of Inorganic Chemistry. 

Antimony Tetroxide : Sb 2 0 4 = (Sb 0 2 ) 0 (SbO), is a white 
powder which becomes yellow when heated, but neither melts nor 
sublimes. It is formed when either antimony trioxide or antimonic 
anhydride is heated to redness in the air, the former compound 
then taking up oxygen and the latter losing it. When fused with 
caustic potash or potassium carbonate, a mixture of potassium 
antimonite and potassium antimonate is formed. 


COMPOUNDS OF ANTIMONY WITH THE 
HALOGENS. 

The haloid compounds of antimony correspond to those of 
phosphorus. Two chlorides are known—viz. the tri- and penta- 
chlorides—both of which, like the chlorides of phosphorus, are de¬ 
composed by water into hydrochloric acid and the corresponding 
oxygen compounds. An antimony compound corresponding to 
phosphoric oxychloride with the composition SbOCl 3 has not yet 
been prepared, but a similar compound of triad antimony--anti- 
monous oxychloride : SbOCl—is known. 


Antimony Trichloride (Antimonous Chloride) : SbCl 3 . 

This compound, sometimes called butter of a?iti7nony from its 
consistency, is a soft white, crystalline solid, which melts at 73 0 , 
boils at 223 0 , and deliquesces in moist air. 

It may be prepared by heating an excess of finely powdered 
antimony in chlorine gas, or by distilling an intimate mixture of 
1 part of powdered antimony with 3 parts of mercuric chloride. 
The usual method of preparation is to dissolve antimony trisulphide 
in strong hydrochloric acid, evaporate the solution to drive off the 
water and excess of acid, and then to distil the residue. As soon 
as the distillate begins to solidify the receiver is changed, and by 
again distilling those portions which come over last, the compound 
is obtained pure. 

Antimony trichloride unites with some metallic chlorides to 
form crystalline double chlorides—^, sodium antimony chloride : 
3NaCI,SbCl 3 . It dissolves in hydrochloric acid or in a small 



Compounds of Antimony with the Halogens. 257 

quantity of water; the addition of a large quantity of water to 
either solution produces a white precipitate of antimony oxy¬ 
chloride (algaroth powder). Tartaric acid prevents this precipi¬ 
tation. 

Algaroth powder is a white amorphous substance insoluble in 
water. It contains antimony, oxygen, and chlorine, but has no 
definite composition. By boiling with water it gives up some 
chlorine, and when boiled with sodium carbonate is converted into 
antimonous acid. 


Antimony Pentachloride (Antimonic Chloride) : SbCl 5 . 

Unlike phosphorus pentachloride, the antimony compound is 
not a solid substance, but a pale yellow liquid, fuming in the air. 
When distilled it undergoes partial decomposition into chlorine and 
antimony trichloride, and does not, therefore, possess any constant 
boiling point. 

Antimony pentachloride is obtained either by saturating the tri¬ 
chloride with chlorine, or by passing a rapid stream of chlorine 
overheated antimony. It attracts moisture from the air, and then 
solidifies to a crystalline mass. In a small quantity of water it 
dissolves to a clear solution, which, on standing over sulphuric acid, 
deposits crystals of a hydrate having the composition SbCl 5 + 
4H 2 0. Excess of water precipitates metantimonic acid, but tartaric 
acid prevents the precipitation. If, however, its acid solution is 
mixed at once with a large quantity of water, the liquid remains 
clear. 

The two atoms of chlorine which antimony pentachloride- con¬ 
tains in excess of the trichloride are only feebly united, and it is 
therefore well adapted for imparting chlorine to many substances 
•— eg. carbonic oxide, ethylene, and other organic compounds. 
With dry sulphuretted hydrogen the pentachloride is converted 
into hydrochloric acid and antimony sulphochloride, SbSCl 3 , a white 
crystalline, easily fusible substance. 


Antimony Tribromide : SbBr 3 , which may be obtained by the 
direct union of its constituents, is a solid crystalline mass, melting 
at 93 0 , boiling at 280° and subliming in colourless needles. The 
pentabromide has not yet been prepared. 


S 


258 Text-Book of Inorganic Chemistry. 

Antimony Triiodide : Sbl 3 , sublimes in large red crystals, 
which become darker when fused. 


SULPHUR COMPOUNDS OF ANTIMONY. 

Two of these compounds are known—viz. the trisulphide and 
the pentasulphide. They correspond in composition to the two 
oxides, and closely resemble one another both in their external and 
in their chemical properties. Both are insoluble in water, and both 
unite with strong bases to form sulphosalts. 

Antimony Trisulpliide (Antimonous Sulphide) ! Sb 2 S 3 , 
occurs in nature as the mineral grey antimony ore , stibnite , or 
antimonite , and is the chief source of antimony and its compounds. 
It is found crystallized in long rhombic prisms or in fibrous crystal¬ 
line masses of a dark grey colour with a metallic lustre. It easily 
fuses and re-solidifies to a crystalline mass, forming the crude im¬ 
pure sulphide of commerce (p. 251). 

It may be artificially prepared by fusing together antimony and 
sulphur, and repeating the process several times with addition of 
sulphur, or by precipitating an acid solution of antimony trichloride 
with sulphuretted hydrogen. The trisulphide prepared in the 
wet way is an amorphous orange-coloured substance, which when 
fused and resolidified closely resembles the ordinary sulphide. 

Antimony trisulphide is converted by nitric acid into insoluble 
antimony trioxide or antimonic acid, while the sulphur partly 
separates in the free state and is partly oxidized to sulphuric acid. 
Hydrochloric acid easily dissolves it, forming antimony trichloride, 
with liberation of sulphuretted hydrogen. It combines easily 
with potassium or sodium sulphide to form a soluble sulphanti- 
monite, from which hydrochloric acid again precipitates the orange 
coloured trisulphide. Caustic soda also dissolves it, and the 
solution then contains sodium antimonite as* well as sodium sulph- 
antimonite : — 

2Sb.,S 3 + 4NaOH = 3SbS-SNa + SbO-ONa + 2H 2 0. 

Sodium Sodium 

sulphantimonite antimonite 




Sulphur Compounds of Antimony. 259 

Not only caustic soda but also sodium carbonate dissolves 
antimony trisulphide, and especially the orange-coloured, amor¬ 
phous variety. The colourless solution when boiled with an excess 
of the trisqlphide takes up still more of it, the greater portion of 
which is again deposited on cooling. This precipitate is not the 
pure sulphide, but contains varying quantities of sodium antimonite 
and antimony trioxide, and was previously used in medicine under 
the name of kermes mineral , or simply kemnes. Ammonium car¬ 
bonate does not dissolve antimony trisulphide, a property which 
serves to distinguish it from the trisulphide of arsenic. 

If antimony trisulphide is heated (roasted) in the air, a por¬ 
tion is converted into sulphurous anhydride and antimony trioxide, 
the latter substance then combining with the undecomposed tri¬ 
sulphide to form a brownish vitreous, semi-transparent mass. 
This mixture of the trioxide and trisulphide, prepared in this and 
other ways, was formerly used in medicine and for the preparation 
of other antimony compounds under the names antimony glass , 
crocus antimonii , &c. It is now only employed for imparting a 
yellow colour to glass and porcelain. 


Antimony Pentasulphide : Sb. 2 S 5 , sometimes called The golden 
sulphide,’ is a dark orange-coloured powder, and may be obtained 
by passing sulphuretted hydrogen through an acid solution of 
antimonic acid. It is, however, usually prepared by precipitating 
a solution of sodium sulphantimonate (Schlippe’s salt) with dilute 
hydrochloric acid. 

Antimony pentasulphide forms with alkaline sulphides a series 
of salts corresponding to a tribasic sulphantimonic' acid. The 
best known and most stable of these is the sodium compound : 
SbS-(SNa) 3 + 9H0O, which is easily soluble in water, and from 
which it crystallizes in colourless tetrahedra. This salt, called 
I Schlippe's salt , after its discoverer, is easily formed by boiling an 
aqueous solution of sodium sulphide with antimony trisulphide, and 
sufficient sulphur to convert it into the pentasulphide. Instead of 
sodium sulphide, caustic soda and sulphur may be used, which on 
boiling yield sodium sulphide and sodium thiosulphate, or, instead 
of caustic soda, sodium carbonate and slaked-lime. To prepare the 
compound 9 parts of crystallized sodium carbonate are boiled with 
3 parts of slaked-lime, 3 parts of antimony trisulphide, 1 part of sul¬ 
phur, and sufficient water. The hot liquid is rapidly filtered off from 




260 


Text-Book of Inorganic Chemistry. 

the calcium carbonate and evaporated down until the salt crys¬ 
tallizes out. The crystal so obtained must be preserved in well • 
stoppered bottles, as the carbonic anhydride of the air decomposes 
them, forming sodium carbonate, sulphuretted hydrogen, and anti¬ 
mony sulphide. This causes the colourless, or, at most, pale yellow 
crystals to become gradually covered with an amorphous brown- 
coloured crust. 

The potassium compound has a similar composition ; aqueous 
solutions of these salts give with most other metallic salts inso¬ 
luble precipitates of corresponding sulphantimonates. The copper 
compound, for example, has the composition : (SbS) 2 (S 2 Cu) 3 . 


BORON. 

Chemical Symbol : B.— Atomic Weight', n. 

Boron, like phosphorus, is only found in inorganic nature in 
combination with oxygen, either as the free compound—boric acid, 
or united with bases as various borates. The most important 
minerals containing boron ar e—sassolite (boric acid) : B(OH) 3 , 
tinkal or native borax : Na 2 B 4 0 7 + ioH 2 0, boracite : 2 Mg 3 B 8 0 15 + 
MgCl 2 , boronatro-calcite (calcium and sodium borates), and datolite 
(calcium borate and silicate). 

Boron can be easily prepared from anhydrous fused boric acid. 
This substance is coarsely powdered, mixed with small pieces of 
sodium (6 parts of sodium to io parts of boric acid), thrown 
into an iron crucible which has been previously heated to bright 
redness, and then covered with 5 parts of well-dried common 
salt. As soon as the reaction in the closed crucible is over its 
contents are stirred with an iron rod and poured, while still fluid 
and red-hot, into dilute hydrochloric acid, in which everything 
dissolves except the boron which has been set free by the sodium. 
The residue is brought on a filter, washed, first with dilute hydro¬ 
chloric acid, then with cold water, and dried at the ordinary tem¬ 
perature on a porous slab, or over sulphuric acid. 

Boron obtained in this way is an amorphous olive-green powder, 
tasteless, and without odour, and a non-conductor of electricity. 
It is extremely infusible. Water, even at its boiling point, does not 



Boron. 


261 


attack it; nor is it acted upon by hydrochloric acid. It is oxidized, 
however, to boric acid by nitric acid, aqua regia, or concentrated 
sulphuric acid, as well as when heated in steam or when fused with 
caustic soda. When heated in the air it burns easily, producing 
boric anhydride; in oxygen it burns with a brilliant light of a 
greenish colour. It also burns in chlorine, producing boron tri¬ 
chloride. Boron unites readily with nitrogen ‘ wheti heated in this 
gas to redness it forms a very stable nitride. 

Boron can also be obtained in a crystalline form by dissolving 
amorphous boron in fused aluminium. Amorphous boron is 
stamped tightly in a Hessian crucible, and a hole made in the 
middle, in which is placed a rod of aluminium. After the crucible 
has been closed, it is placed inside a second large crucible, and 
the space between them filled up with powdered charcoal; the lid 
of the exterior crucible is then cemented with clay, and the whole 
heated to bright redness for two hours. On cooling, the surface of 
the aluminium is found covered with crystals of boron. The 
aluminium is dissolved by dilute hydrochloric acid, and the crys¬ 
talline boron then remains, partly as dark brown translucent, 
and partly as transparent yellowish crystals of specific gravity 
= 2-6, mixed with thin, opaque, six-sided, tabular crystals of a 
compound of aluminium and boron, which may be easily removed 
by washing. 

Another kind of crystalline boron distinguished by its great 
hardness is obtained by strongly heating a mixture of boric anhy¬ 
dride and aluminium in a carbon crucible, air being excluded as 
much as possible. These crystals, which approach the diamond in 
brilliancy and hardness, are not, however, pure boron, but contain 
4 per cent, of carbon and 67 per cent, of aluminium. 

No practical use has yet been discovered for boron. A com¬ 
pound of boron with hydrogen is unknown, and in general the 
number of the boron compounds is small. By far the most impor¬ 
tant compound is boric acid, in which, as in all other compounds, 
boron plays the part of a triad element. 






262 


Text-Book of Inorganic Chemistry . 


BORIC ANHYDRIDE : B 2 0 3 , and BORIC ACID : B(OH) s . 

Boric acid crystallizes in colourless lustrous tablets, with a 
faint acid taste. It is soluble in water, especially when warm, and 
also dissolves in alcohol. One part of boric acid dissolves in three 
parts of water at ioo°, but requires twenty-five parts at 19 0 . A hot 
saturated solution deposits therefore nearly all the boric acid on 
cooling. The aqueous solution possesses both a faint acid and an 
alkaline reaction; it reddens blue litmus paper, and at the same 
time turns yellow turmeric paper brown. 

Boric acid occurs in such large quantities free in nature that, 
until quite recently, other naturally occurring boron compounds 
have been but little used for the preparation of the acid and its 
salts. In some parts of Tuscany jets of steam (called suffio 7 ii , or 
fumaroli ) issue from the ground, which then condense to form 
bogs or marshes {lagoons). This steam carries with it small quan¬ 
tities of boric acid, and although the condensed water rarely con¬ 
tains more than ~ 6 per cent, of the compound, many hundred¬ 
weights of boric acid are annually prepared from this source. 

The places where these jets of steam issue from the earth, or 
where they have been artificially produced by boring, are built 
round so as to form a large basin, which is then filled with water. 
The water is soon raised to boiling by the condensed steam, from 
which it abstracts the boric acid. After a certain time the water 
is allowed to flow into a second somewhat lower basin where it re¬ 
ceives fresh quantities of boric acid, while the first basin 'is again 
filled with fresh water. When the water has, in this way, passed 
through four or more basins, from each of which it has received 
boric acid so that it contains about one per cent., the suspended 
impurities are allowed to settle and the liquid evaporated down to 
crystallize. For this purpose it is run into shallow pans heated 
from beneath by other similar jets of steam. And in this way the 
use of fuel, which it is difficult to procure in the district, is entirely 
avoided. The boric acid which separates out on cooling the solu¬ 
tion. when sufficiently concentrated, is contaminated with various 
impurities, especially with ammonium and calcium sulphates. It 
is purified by recrystallization. 

Similar steam-springs containing boric acid are also met with 
in other volcanic districts, for example in the Lipari Islands. 


Boric Anhydride and Boric Acid. 263 

Recently considerable quantities of boric acid have been dis¬ 
covered in California, whence most of the American acid is now 
derived. 

Several hypotheses have been proposed to explain the occurrence 
of boric acid in these hot springs ; it is most probable that the acid 
is produced by the action of steam on boron nitride, which is then 
decomposed into boric acid and ammonia. Ammonium com¬ 
pounds are always present in the hot springs. 

The extraction of boric acid from native borax (tinkal) and from 
other boron minerals, especially boronatro-calcite and borocalcite, 
consists simply in decomposing them with hot hydrochloric acid 
and purifying the boric acid which separates out on cooling by re¬ 
crystallization. 

Although boric acid is only slightly volatile even at high tem¬ 
peratures, we have seen that it passes off with the steam when its 
aqueous solution is boiled. Even when its alcoholic solution is 
boiled considerable quantities are carried off with the alcohol 
vapour, which then burns with a bi'ight green flame when ignited. 
This colour is, however, only imparted to the flame of the burning 
alcohol by free boric acid, not by its salts. Alcohol, if poured on 
sodium borate or any other salt of boric acid and ignited, burns 
with a yellow flame ; but the addition of a drop of concentrated 
sulphuric acid sets the boric acid free, and the flame changes at 
once to a bright green. 

When heated to ioo° ordinary boric acid : B(OH) 3 , loses water 
and changes into metaboric acid, a white powder of the composition : 
BO • OH ; by heating to 150°—160° in a stream of dry air a further 
quantity of water is lost and an acid of the composition : H 2 B 4 0 7 , 
pyroboric acid , remains. Heated still more strongly, boric acid 
froths up, and at red-heat melts to a clear viscid liquid, consisting 
of boric anhydride, which solidifies to a hard transparent brittle 
glass. On standing in the air, boric anhydride gradually becomes 
opaque owing to the absorption of water, for which it possesses 
considerable attraction. 

Boric acid is a~weak acid, as is shown by its behaviour with 
litmus and turmeric. It combines easily with the alkalies to form 
salts soluble in water; the other borates are difficultly soluble in 
water, but easily soluble in acids. 

The constitution of the borates is usually somewhat complex. 
Compounds are known of the ordinary tri-basic acid, B(OH) 3 , par¬ 
ticularly organic salts, in which the three hydrogen atoms are dis- 


264 Text-Book of Inorganic Chemistry. 

placed by compound radicals. Salts have also been prepared from 
the mono-basic metaboric acid : BO-OH, in which the one atom of 
hydrogen is displaced by an atom of a monad metal. Most of the 
borates, including the commonest—sodium borate or borax—have, 
however, a more complex composition, and correspond to the di¬ 
basic pyroboric acid, H 2 B 4 0 7 — eg. Na 2 B 4 0 7 , and CaB 4 0 7 . 

The constitution of these pyroborates may be represented in 
various ways. Formerly it was considered that borax, which may 
be taken as representing this class, was a double compound of 
sodium metaborate with boric anhydride: 2BO • ONa + B. } 0 3 . A 

more probable formula is : O j^'o'-ONa’ or ’ P ossibl y> a portion 

of the boron may be present as acid, and a portion as base; but 
until our knowledge is further advanced no constitutional formula 
can be given with certainty. 

Boric acid, like all weak acids, can unite with strong acids, and 
forms compounds with sulphuric and phosphoric anhydrides having 
the respective composition : B 2 0 3 ,S 0 3 and B 2 0 3 ,P 2 0 5 , in which 
the boric acid plays the part of a base. 

The chief use of boric acid is for the preparation of borax. It 
is an excellent antiseptic, and has been considerably used for this 
purpose in recent years. Whether, however, large quantities of 
boric acid can be taken into the system without injurious effects 
remains as yet doubtful. 


OTHER COMPOUNDS OF BORON. 

Boron Trichloride : BC 1 3 .—This compound is produced when 
boron is heated in gaseous chlorine, or by strongly glowing an 
intimate mixture of boric anhydride and charcoal in a porcelain 
tube through which a stream of dry chlorine is passed. The 
reaction which then goes on is expressed by the following equa¬ 
tion :— 

BA + 3 C + 3 C 1 2 = 2 BC 1 3 + 3 CO. 

The gaseous products are led through a U-tube surrounded by 
a freezing mixture, in which the boron trichloride condenses to 
form a.colourless.liquid, boiling at 17 0 , and having a specific gravity 



Other Compounds of Boro?i. 265 

of 1-35. This liquid fumes strongly in the air, and is rapidly de¬ 
composed by water into boric and hydrochloric acids. Boron 
trichloride unites with ammonia to form a solid white crystalline 
compound of the composition : 2BC1 3 ,3NH 3 . This compound, 
when heated, sublimes, and is decomposed at once by water into 
boric acid, hydrochloric acid, and ammonium chloride. 

Boron Tribromide : BBr 3 , is a colourless liquid, boiling at 90°, 

1 and with similar properties to the chloride. 

Boron Trifluoride : BF 3 .—This body, unlike the two other 
j corresponding compounds, is a colourless gas, and can only be 
condensed to a liquid with difficulty. It possesses a piercing 
| suffocating odour, fumes strongly in the air, and on account of 
its strong attraction for water, chars many organic substances like 
concentrated sulphuric acid. It is extremely soluble in water, one 
j volume of this liquid dissolving 700 to 1,000 volumes of the gas. 

I The gas may be obtained by heating an intimate mixture of 1 part 
I of boric anhydride (or 2 parts of anhydrous borax) with 2 parts of 
fluor-spar and 12 parts of concentrated sulphuric acid in a glass 
flask. The change which then takes place is as follows :— 

B 2 0 3 + 3 CaF 2 + 3 S 0 2 (OH) 2 = 3 S 0 2 . 0 2 Ca + 2 BF 3 + 3 H 2 0 . 

The large excess of sulphuric acid is required to absorb the 
water produced in the reaction. Boron trifluoride may also be 
prepared from boric anhydride and fluor-spar without sulphuric 
acid if a mixture of these two substances is heated to bright red¬ 
ness in an iron tube, the calcium then remaining behind as calcium 
borate : — 

2 B 2 0 3 + 3 CaF 2 = 2BF 3 + B 2 ( 0 2 Ca) 3 . 

On evaporating an aqueous solution of the fluoride, boric acid 
j separates out, and a monobasic acid remains, of the composition : 

I HBF 4 or HF,B ¥ z —fuoboric acid— which may be considered as 
I a double compound of hydrofluoric acid and boric fluoride. The 
same substance is also more easily obtained by dissolving boric 
I acid in hydrofluoric acid. 

Fluoboric acid cannot, however, be obtained in the pure state. 
If the aqueous solution is concentrated by evaporation, hydro- 
' fluoric acid is evolved. Its salts, having the general composition : 
M'BF 4 (where M' is a monad metal), are mostly soluble in water, 
the most insoluble being the potassium compound : KBF 4 . On 



266 Text-Book of Inorganic Chemistry. 

mixing aqueous fluoboric acid with a solution of a potassium 
compound, a gelatinous precipitate of potassium fluoborate is 
formed, which becomes a white powder when dried. From a hot 
saturated solution in water, the salt separates out on cooling in 
brilliant crystals. 

Boron Sulphide : B. 2 S 3 , is a white, vitreous solid, of piercing 
odour, which can be fused in a stream of hydrogen. It is pro¬ 
duced by heating amorphous boron in sulphur vapour or in sul¬ 
phuretted hydrogen, and then condenses in a well-cooled receiver. 
Water easily decomposes it into boric acid and sulphuretted 
hydrogen. 

Boron Nitride : BN.—Boron is one of the few elements that 
unite directly with nitrogen. The nitride is formed when amor¬ 
phous boron is heated in a stream of nitrogen or ammonia, also 
when a mixture of boric anhydride and charcoal is strongly heated 
in a stream of nitrogen, or by glowing a mixture of dried borax and 
ammonium chloride. By digesting the residue obtained in the last 
method with dilute hydrochloric acid, the nitride remains as a 
white amorphous powder. Boron nitride is a stable compound, 
and remains unchanged when boiled with water or when glowed 
in the air or in hydrogen. It is, however, decomposed into boric 
acid and ammonia when heated to low redness in a current of 
steam ; the same change is produced by fusion with caustic potash. 


SILICON. 

Chemical Symbol : Si .—Atomic Weight : 28. 

Silicon occurs most abundantly and is widely distributed in 
nature, but never in the free state. It is always found, like phos¬ 
phorus and boron, united with oxygen. Its only oxide, silica, is known 
both free and in combination with bases as innumerable minerals. 
Considering the wide distribution and great variety of the silicates, 
it may be asserted that silicon plays the same part in inorganic as 
carbon does in organic nature. Both are tetrad elements in almost 
all their compounds. 



Silicon. 


267 

Silicon may be separated from silica by potassium, but the 
decomposition is always incomplete, and the product therefore 
impure. It is better to employ a compound of silicon fluoride with 
potassium or sodium fluoride—potassium or sodium fluosilicate : 
K 2 SiF 6 —which can be readily obtained in the pure state. A 
mixture of one of these salts with sodium chloride and metallic 
sodium cut into small pieces is thrown into a red-hot iron crucible, 
which is then closed and kept for a short time at a low red heat. 
After cooling, the contents of the crucible are boiled with dilute 
hydrochloric acid. By this process fluorine is abstracted by the 
sodium from the double fluoride and silicon set free. 

Amorphous silicon prepared in this way is a dark brown powder, 
insoluble in water and in nitric or sulphuric acid. Aqueous hydro¬ 
fluoric acid dissolves it with evolution of hydrogen, and it also de¬ 
composes strong caustic potash, forming potassium silicate. Heated 
in the air it burns easily with a bright light, to form silica ; the com¬ 
bustion, however, is only incomplete, as the silica is fused by the high 
temperature, and so protects a portion of the silicon from the air. 
When heated in a stream of hydrochloric acid gas this form of silicon 
takes up three atoms of chlorine and one of hydrogen, and becomes 
converted into the substance called silicochloroform : SiHCl 3 (see 
sequel). If amorphous silicon is strongly heated in a crucible out 
of access of air, it contracts considerably, becomes of a chocolate 
brown colour, and now no longer catches fire when heated in the 
air. It has also lost its property of dissolving in hydrofluoric acid 
or in caustic potash after this treatment. 

Silicon may be prepared in the crystalline form by fusing 
aluminium with thirty times as much sodium fluosilicate. The 
black residue which remains is then treated first with concentrated 
hydrochloric acid to remove the aluminium, and then with hydro¬ 
fluoric acid. Another and simpler method of preparing crystalline 
silicon consists in throwing a mixture of three parts of potassium 
fluosilicate, one part of sodium in small pieces, and one part of 
granulated zinc, into a red-hot Hessian crucible, and then heating 
for some time to such a temperature that the zinc remains fused, 
but is not volatilized. The zinc regulus, which contains the silicon, 
is afterwards removed from the crucible, thoroughly boiled with 
water to remove the slag, and then the zinc dissolved out with 
hydrochloric acid. 

Crystalline silicon prepared in either of these ways consists 
either of opaque, lustrous tablets of a dark grey colour, resembling 


268 Text-Book of Inorganic Chemistry. 

graphite, or else of brilliant prismatic crystals of an iron grey 
colour and considerable hardness. The specific gravity of this 
modification of silicon is 2*5 ; it is a conductor of electricity, and 
remains unaltered when heated in the air, but melts at high tem¬ 
peratures. It is neither attacked by nitric nor by hydrofluoric acid, 
but dissolves when heated with strong caustic potash, with evolu¬ 
tion of hydrogen, and burns in chlorine to silicon tetrachloride. 


COMPOUNDS OF SILICON. 

Silicon unites directly with oxygen and with the halogens to 
form compounds, in which it exists as a tetrad element. An oxide 
of silicon corresponding to carbonic oxide, with the composition : 
SiO, has not yet been prepared. Otherwise its compounds are 
similar to those of carbon, although the similarity is more in their 
composition than in their properties. The following pairs of com¬ 
pounds correspond to one another 


Silica 

. Si 0 2 

Carbonic acid . 

. C 0 2 

Silicon hydride 

. SiH 4 

Methane (marsh-gas) 

. CH 4 

Silicon tetrachloride 

. SiCl 4 

Carbon tetrachloride 

. CC 1 4 

Silicon hexachloride 

. Si 2 Cl 6 

Carbon hexachloride 

• C 2 C 1 6 

Silicon sulphide 

. SiS 2 

Carbon sulphide 

. cs 0 

Silicochloroform 

SiHClj 

Chloroform 

chc£ 

Silicoiodoform 

. SiHI 3 

Iodoform. 

. CHI3 

Silicoxalic acid 

(SiO-OH 
| SiO -OH 

Oxalic acid 

jCO-OH 

iCO-OH 


SILICA (SILICIC ANHYDRIDE): Si 0 2 , 
and SILICIC ACID. 

Silicon forms only one oxide—silicic anhydride—or, as it is 
briefly called, silica. This united in varying proportions with water 
forms different varieties of silicic acid. 

Free silica occurs in nature both crystalline and amorphous. 
In the crystalline form chiefly as quartz or rock crystal, in six- 



Silicic Acid . 


269 

sided prisms bounded at the ends by similar pyramids, and belong¬ 
ing therefore to the hexagonal system. Quartz in its purest form 
is transparent and colourless, but is frequently coloured brownish or 
violet—in the former case it is known as smoky-quartz, in the latter 
as amethyst. Ordinary sand consists of particles of silica, which 
when united together by some cementing substance constitute the 
different varieties of sandstone. Quartz is also contained in the 
free state in many important rocks — granite , for example, is made 
up of separate crystals of quartz, felspar, and mica. The specific 
gravity of this form of silica is 2-65. 

Crystalline silica is also found in some rocks in minute crystals 
belonging to the rhombic system, and has a specific gravity of 2’31, 
or lower than that of quartz. This variety, which is called tridy- 
mite , appears to be the most stable form of silica at a red-heat, as 
both quartz and amorphous silica pass into it when strongly heated. 

Besides these two crystalline forms, silica is also found in nature, 
usually in combination with water, in the amorphous state. The 
minerals agate, chalcedony, opal, jasper, flint, consist of amorphous 
silica mixed to some extent with the crystalline modification. 
Spme of these— e.g. opal—are of attractive colours and are used as 
jewels. Agate and chalcedony are celebrated for their extreme 
hardness, and are highly valued for the manufacture of mortars. 
The siliceous sinter, deposited during the evaporation of the water 
from siliceous springs, and the so-called infusorial earth, or the in¬ 
organic remains of certain infusoria, consist essentially of amor¬ 
phous silica. This infusorial earth— Kieselguhr of the Germans— 
occurs largely in certain localities in North Germany as a light, 
extremely fine powder, mostly of a yellow colour. It is used for 
many technical purposes, especially for mixing with nitro-glycerine 
to form dynamite. 

Silica is likewise found in the animal and vegetable kingdoms. 
The glassy coating of the stems of certain grasses and other allied 
plants, cereals, rushes, and especially of the Equisetaceae (horse¬ 
tails) consist essentially of silica. The ashes of the feathers of many 
birds contain as much as 40 per cent, of silica. 

Much more widely distributed in the mineral kingdom than 
silica are its compounds with various bases, called silicates, and of 
which the greater part of the solid crust of the earth consists. 
Among these the different felspars, double silicates of aluminium 
and potassium or sodium, together with the clays, schists, &c., 
produced from them, take the first place. 


270 Text-Book of Inorganic Chemistry. 

The properties of the various modifications of silica differ in 
many respects. All are insoluble in water and in nitric, hydro¬ 
chloric, or sulphuric acid, as well as in aqua regia, but dissolve in 
hydrofluoric acid, and possess in general none of the properties of 
an acid. They remain unchanged at the highest temperatures of 
our furnaces, but melt in the oxy-hydrogen flame and form a clear 
transparent glass on cooling. The quartz modification is very 
hard—it easily scratches glass—but the amorphous form is much 
softer. The specific gravity of the former is 2*65, that of the latter 
2-20 ; the specific gravity of both changing to 2-31 when strongly 
glowed, owing to the conversion into tridymite. Crystalline silica 
is scarcely attacked even by boiling aqueous caustic potash 
or soda, but when fused with an excess of either of these 
substances in the solid form, or with their carbonates, silicates 
soluble in water are produced (water-glass). Amorphous silica is 
dissolved more or less readily by aqueous alkalies ; infusorial earth 
is very readily dissolved. 

If a concentrated aqueous solution of an alkaline silicate pre¬ 
pared by either of the above-mentioned methods is mixed with 
hydrochloric acid, the silicic acid separates out as a gelatinous 
mass, and the whole liquid becomes so viscid that the vessel may 
be inverted without anything running out. This silicic acid is 
soluble in water, especially in the presence of hydrochloric acid, 
for if the solution of the silicate is sufficiently dilute, hydrochloric 
acid produces no precipitate. It is this silicic acid which is con¬ 
tained in the water of siliceous springs, and which is deposited as 
siliceous sinter on evaporation. 

Pure amorphous silica, as a light, finely divided powder, may be 
easily obtained by the decomposition of silicon fluoride with water 
(p. 275).. 

If this gelatinous soluble silicic acid is evaporated to complete 
dryness on a water-bath, it loses its water, and there remains 
amorphous silica, which is no longer soluble in water or hydro¬ 
chloric acid, but which easily dissolves on warming in aqueous 
caustic soda or in sodium carbonate. When heated to redness it 
is no longer soluble in aqueous alkalies, and can now only be con¬ 
verted into soluble water-glass by fusion with caustic alkalies or 
alkaline carbonates. 

Corresponding to these differences in solubility of the various 
forms of silica are differences in the readiness with which the 
various silicates are decomposed. Those which are poor in silica, 


Silicic Acid. 


271 


especially those containing water (the zeolites), when finely divided, 
are decomposed by hydrochloric acid. Those, on the other hand, 
which contain more silica—^, the felspars—withstand the action 
of strong hydrochloric acid, and can be only brought into solution 
by treatment with hydrofluoric acid or by fusion with sodium 
(potassium) carbonate in a platinum crucible. In the latter case 
an alkaline silicate soluble in water is produced, and the bases 
which were previously united with the silica— e.g. lime, alumina, 
iron—are converted either into carbonates or into oxides soluble in 
acids. The fused cold mass is now digested with water and 
hydrochloric acid added to slight excess, when a clear solu¬ 
tion is obtained as soon as the carbonic acid has been expelled 
by warming. If the quantity of water used was small, the solution 
will contain gelatinous flocks of silicic acid. The liquid is next 
evaporated to complete dryness, and best heated for a short time 
to about 120 0 , to render the silica completely insoluble, afterwards 
moistened with strong hydrochloric acid to bring into solution any 
oxychlorides which might have been formed during evaporation 
(e.g. of iron or aluminium), and finally extracted with hot water. 
In this way everything goes into solution except the silica, 
which can be easily separated by filtration, washed and weighed, 
while the bases in the filtrate can be estimated by any known 
method. 

It is difficult to obtain a silicic acid of definite composition. If 
the gelatinous hydrate is dried over concentrated sulphuric acid, a 
transparent vitreous mass is obtained, of which the composition is 
approximately represented by the formula : SiO(OH) 2 —z>. a dibasic 
acid corresponding to the hypothetical sulphurous acid. Tetra- 
basic silicic acid of the composition : Si(OH) 4 has not yet been pre¬ 
pared. We know, however, from the composition of numerous 
silicates that salts of both these acids exist, as well as compounds 

of a dibasic disilicic acid of the possible formula : O j g|o -OH 

To the salts of the dibasic acid may possibly belong the mineral 
wollastonite : Si 0 - 0 2 Ca, to those of the tetrabasic acid, olivine : 
Si (0 Mg) 2 . And a combination of the dibasic and tetrabasic acids 

may represent serpentine : g] I ( 0 2 Mg) 3 . 

Similarly, those silicates containing three atoms of silicon—the 
trisilicates—may be considered as compounds of two of these acids 


272 


Text-Book of Inorganic Chemistry. 


united together. Orihoclase (potash felspar : KAlSi. H 0 8 ) would 
thus be a compound of dibasic silicic acid and disilicic acid :— 



and pretuiite : AUCa.^SigOn, a compound of tetrabasic silicic acid 
and of a hexabasic disilicic acid of the hypothetical constitution : 



It must, however, be distinctly understood that these formulas 
are not the symbolic expression of established facts, but only re¬ 
present an attempt to give a simple explanation of the very various 
and often complex composition of the silicates. The chemical 
cojistitution of these compounds cannot be deduced simply from 
the results of analysis, but require extended researches similar 
to those which are necessary to Establish the constitution of any 
chemical compound. 


SILICON HYDRIDE. 


Composition : SiH 4 . 


This compound is a colourless combustible gas, insoluble (or 
nearly so) in water. It may be obtained by acting upon magne¬ 
sium silicide with dilute hydrochloric acid :— 

SiMg 2 + 4HCI = 2MgCl., + SiH 4 . 

This magnesium silicide (SiMg,,) is prepared by throwing a 
mixture of 8 parts of anhydrous magnesium chloride, 7 parts 
of sodium fluosilicate, 2 parts of sodium chloride, and 4 parts 
of sodium in small pieces into a red-hot Hessian crucible, which 
is at once closed. The crude compound thus obtained always 
contains free magnesium, whence it follows that the silicon 
hydride prepared from it is always mixed with free hydrogen. In 
this state silicon hydride catches fire spontaneously in the air, 



Silicon Tetrachloride. 273 

burning with a luminous flame to form silica and water. If a clean 
porcelain dish is depressed on the flame, brown amorphous silicon 
is deposited on it. Excluded from the air and heated to faint 
redness, the gas decomposes, like arsine and stibine, into amorphous 
silicon and free hydrogen. 

Th zpure compound may be prepared by acting on an organic 
silicon compound of the composition : SiH(OC 2 H 5 ) 3 , with sodium. 
Four molecules of this compound are decomposed under the 
influence of the sodium, which remains itself unchanged, into 
one molecule of silicon hydride and three of the compound • 
Si(OC 2 H 5 ) 4 

4SiH(OC 2 H 5 ) 3 = SiH 4 + 3Si(OC 2 H 5 ) 4 . 

Silicon hydride burns easily in chlorine, producing hydrochloric 
acid and silicon tetrachloride. With caustic potash, it gives a 
silicate of potassium and free hydrogen :— 

SiH 4 + 2KOH + H 2 0 = SiO(OK) 2 + 4 H r 


SILICON TETRACHLORIDE. 

Composition : SiCl 4 . 

This, the best known chloride of silicon, is a colourless mobile 
liquid of 1*52 specific gravity and boiling at 58°. It fumes in the 
air, and is energetically decomposed by water into hydrochloric 
and silicic acids. 

Silicon tetrachloride is obtained when amorphous silicon is 
heated in a stream of chlorine gas, or more readily by strongly heat¬ 
ing an intimate mixture of silica and carbon in the same gas. A 
stiff dough is made of finely divided silica, powdered wood-char- 
coal, and starch paste, which is then formed into balls. These are 
allowed to dry in the air, and then heated to bright redness 
imbedded in charcoal powder. The porous mixture of silica and 
carbon so obtained is introduced while still hot into a dry porcelain 
tube, which can be heated to bright redness in a tube-furnace, 
while a current of dry chlorine passes through it. The reaction is 
expressed by the equation 

Si 0 2 + 2C + 2C1 3 = SiCl 4 + 2CO. 

The other end of the tube communicates air-tight with a 

T 



274 Text-Book of Inorganic Chemistry. 

U-tube placed in a freezing mixture in which the silicon tetra¬ 
chloride condenses, while the carbonic oxide and excess of chlorine 
are carried away to the open air. At the bottom of the U-tube is 
a small vertical tube (provided with a glass stopcock), through 
which the silicon tetrachloride can be drawn off into bottles or 
sealed tubes. 

The corresponding silicon tetrabromide , SiBr 4 , and the ietno- 
dide, Sil 4 , are prepared in exactly the same way as the chloride. 
The bromide is a colourless liquid, boiling at 153 0 and solidifying 
to a crystalline mass at -12 0 . The iodide is a crystalline solid, 
boiling at about 290°. 

Silicon Hexachloride : Si 2 Cl 6 .—Besides the above-mentioned 
tetrachloride of silicon, a second compound—silicon hexachloride : 
Si 3 Cl 6 —is also known. This is also a colourless liquid, but has a 
specific gravity of 1*58, boils at 145 °? an< ^ solidifies at about —4 • 
It is obtained by heating silicon in the vapour of silicon tetra¬ 
chloride or by heating a mixture of silicon tri-iodide and mercuric 
chloride. At 350° it decomposes into the tetrachloride and free 
silicon, and this decomposition increases up to 8oo°. But at about 
i,ooo° the hexachloride is again stable— i.c. silicon and the tetra¬ 
chloride again unite to form the hexachloride. Water decomposes it 

into hydrochloric acid and disilicic or silico-oxalic acid : |giO-OH 

The corresponding silicon hexiodide is a crystalline solid, ob¬ 
tained by heating the tetriodide with finely divided silver. Silver 
iodide is then formed with the hexiodide. 


SILICON FLUORIDE. 

Composition : SiF 4 . 

This substance is a colourless gas, fuming strongly in the air, 
and with a suffocating odour. Its specific gravity is 3*6, and it is 
at once decomposed by water with separation of gelatinous silicic 
acid. It is produced when hydrofluoric acid acts upon silica in the 
presence of de-hydrating substances. 

For its preparation a mixture of powdered fluor-spar and fine 



Fluosilicic Acid. 


275 

dry sand, the latter in considerable excess (about equal parts), is 
heated in a thick glass flask with a large excess of concentrated 
sulphuric acid, so that a thin liquid of the consistency of cream is 
produced. Silicon fluoride is then evolved, according to the equa¬ 
tion :— 

Si 0 2 + 2CaF 2 + 2 S 0 2 ( 0 H) 2 = SiF 4 + 2 S 0 3 ( 0 2 Ca) + 2 H 2 0 , 

while the water produced at the same time is absorbed by the 
excess of sulphuric acid. If the quantity of sand employed in the 
experiment is insufficient, the glass flask is strongly attacked by 
the hydrofluoric acid, and may be completely eaten through. 


FLUOSILICIC ACID. 

Co 7 nposition : H 2 SiF 6 . 

This dibasic acid, which maybe considered as a double fluoride 
of silicon and hydrogen, is produced by the union of silicon fluoride 
with nascent hydrofluoric acid. These conditions are fulfilled when 
silicon fluoride comes into contact with water. Decomposition 
then occurs according to the equation :— 

SiF 4 + 3 H 2 0 = SiO(OH) 2 + 4 HF. 

or, 

SiF 4 + 4 H 2 0 = Si(OH) 4 + 4 HF. 

The silicic acid separates as a gelatinous mass, and the hydro¬ 
fluoric acid unites at once with another portion of silicon fluoride, 
forming fluosilicic acid, which remains in solution and can after¬ 
wards be separated by filtration. The whole reaction may, there¬ 
fore, be thus represented :— 


3 SiF 4 

+ 3H 2 0 

= SiO(OH) 2 

+ 2H 2 SiF 6 , 

3 SiF 4 

+ 4 h 2 o 

= Si(OH) 4 

+ 2H 2 SiF 6 . 


The glass tube leading the silicon fluoride into the water, even 
when wide, becomes easily stopped up by the deposited silicic acid, 
and so an explosion might be produced. In order to avoid such 
an occurrence a little mercury is placed at the bottom of the vessel 
in which the fluosilicic acid is to be produced. The tube from the 
flask evolving the silicon fluoride is made to dip under this mercury, 

T 2 



276 Text-Book of Inorganic Chemistry. 

and the vessel then filled up with water. The decomposition of 
the silicon fluoride then only commences above the level of the 
mercury. The danger of an explosion may also be avoided by 
connecting the end of the delivery tube with the neck of a glass 
funnel, and dipping the large open end in the water, so that the 
gas comes at once into contact with a large surface of water. 

The aqueous fluosilicic acid, when separated from the silicic 
acid by filtration, is a strongly acid colourless liquid. It may,be 
concentrated up to a certain point by evaporation at the ordinary 
temperature of the air. But if this limit is exceeded, or if the 
liquid is heated, it decomposes into silicon fluoride and hydro¬ 
fluoric acid, both of which volatilize. On heating, therefore, a 
solution of fluosilicic acid in a platinum basin, ultimately nothing 
remains. 

By neutralizing fluosilicic acid with bases or carbonates, its 
salts—the fluosilicates—are obtained, having the general composi¬ 
tion : R' 2 SiF 6 , where R/ is one atom of a monad metal. These 
salts are 'mostly soluble in water, except the barium compound : 
BaSiF 6 , and those of potassium and sodium. As the strontium 
salt is easily soluble in water, fluosilicic acid has been used as a 
means of separating barium and strontium. 


In concluding this account of the compounds of silicon, brief 
reference may be made to certain substances which are of interest 
on account of analogy with similar compounds of carbon usually 
included under the head of organic chemistry (compare p. 268). In 
lack of suitable names for these compounds they are usually called 
after the corresponding carbon compounds; thus the body corre¬ 
sponding to chloroform (CHC 1 3 ) is called silico-chloroform. 


Silico-chloroform : SiHCl 3 , is produced when amorphous 
silicon is heated in a stream of dry hydrochloric acid gas. At the 
same time a small quantity of the less volatile tetrachloride is also 
produced, which can be separated by fractional distillation. This 
compound is a colourless inflammable liquid, fuming in the air, with 
a specific gravity of 1-65 and boiling at 36°. Chlorine converts it into 
silicon tetrachloride and hydrochloric acid, and water, even at o°,into 
hydrochloric acid and a compound of the composition : Si 2 H 2 0 3 , 


HCO 1 


corresponding to the hypothetical formic anhydride : j O. 


Carbon. 


2 77 


Silico-formic Anhydride : pjgjQ j O, produced as just men¬ 
tioned, is a white powder, which when heated in the air burns easily 
to silica and water. Caustic soda easily dissolves it, with evolution 
of hydrogen, forming sodium silicate. 


Silicoxalic Acid : jsjQ.Qjj 1S obtained as a white powder by 

decomposing silicon hexachloride with water, and is dissolved by 
caustic soda to form sodium silicate. In composition it corresponds 


( CO • OH 

to the carbon compound, oxalic acid: |(^q*oh 

however, it has no other similarity; it does not even possess the 
properties of an acid. 


CARBON. 

Chemical Symbol : C .—Atomic Weight : 12. 

Carbon occupies a peculiar position among the other elements. 
Although it only combines directly with few other elements— 
with oxygen, sulphur, hydrogen, and iron, but neither with nitrogen 
nor the halogens—the compounds of carbon far exceed those of 
any other element both in number and variety. All organic com¬ 
pounds, including the constituents and products of all animal and 
vegetable organisms, as well as the numerous substances which 
can be artificially prepared from them, are compounds of carbon. 

Organic chemistry has been defined as the chemistry of the 
carbon compounds, according to which a description of these sub¬ 
stances should find no place in a text-book of inorganic chemistry. 
But sharp lines of demarcation, which so completely separate sub¬ 
stances from one another as to make a strict classification possible, 
are unknown in nature. Connecting links are eveiywhere to be 
found showing a gradual transition from one class of substances to 
the other. And although all organic compounds contain carbon, 
this by no means compels us to refer all carbon compounds to 
organic chemistry; many of the carbon compounds, especially 
those of simpler composition, cannot be easily separated from the 
inorganic compounds of other elements. 

Carbon occurs free in nature in three modifications, as different 



2/8 Text-Book of Inorganic Chemistry. 

from one another in their physical properties as could be imagined. 
We know it, firstly, as the diamond, crystalline, transparent, cha¬ 
racterized by its great hardness and high index of refraction for 
light, and a non-conductor of electricity; secondly, as graphite, 
with a dark grey colour, crystalline, lustrous, a conductor of elec¬ 
tricity, and extremely soft ; and, lastly, as charcoal, the chief con¬ 
stituent of common coal, amorphous, black, not much harder than 
graphite, and a non-conductor of electricity. 

The diamond is among the rarest of substances and is the most 
precious of gems ; graphite is found in large quantities, but usually 
in isolated spots ; amorphous carbon, as the various forms of coal, 
occurs widely distributed and in immense quantities. The last- 
named variety, although by far the cheapest, is still by far the most 
valuable and important. It furnishes the most important source of 
our artificial heat, light, and power, and influences not only the 
industrial, but also the entire social and political life of the civilized 
countries in which it occurs. 

Like their other physical properties, the specific gravity of the 
three modifications of carbon is very different. That of the 
diamond is 3-5, of graphite 2-2, and of amorphous carbon i*6 
to 2*0. 

Greater similarity, however, exists between the chemical pro¬ 
perties of the three forms. Although they burn in oxygen or air 
with different degrees of readiness, they all produce the same pro¬ 
ducts of combustion. Equal, weights of either modification always 
produce equal weights of carbonic acid. And by the reduction of 
carbonic acid, carbon is again formed, but only amorphous carbon 
or graphite, never diamond. All attempts to prepare the diamond 
by the reduction of carbonic acid or by any other method have, as 
yet, been unsuccessful. 

Carbon in all three forms is infusible at all temperatures, but 
volatilizes at the high temperature of the electric arc ; it is insoluble 
in any of our ordinary solvents. The only solvent at present known 
is fused iron, which, however, only dissolves 1 or 2 per cent, of 
it. On cooling, most of the carbon separates out in the form of 
graphite. Should a liquid ever be discovered capable of dissolving 
carbon at not too high a temperature, the carbon might be again 
separated in the diamond form. 

As was stated above, carbon does not unite directly with nitro¬ 
gen nor with the halogens, and carbon compounds of these sub¬ 
stances can only be obtained indirectly. Even with oxygen, 


Carbon. 


279 


sulphur, and hydrogen, carbon has not the slightest tendency to 
combine, either at the ordinary temperature or at ioo°. But when 
heated to redness its affinity for oxygen and sulphur gradually in¬ 
creases, so that at a white heat the affinity of carbon for oxygen 
and sulphur even surpasses that of phosphorus and potassium. 
Carbon when heated to redness is, therefore, one of the most 
powerfully reducing substances known, and a very extended use 
is made of this property in the extraction of metals from their 
ores. 

Hydrogen only unites with carbon at the extremely high tempe¬ 
rature of the electric arc. When this is produced between carbon 
poles in an atmosphere of hydrogen, small quantities of a com¬ 
pound of the two bodies, called acetylene : C 2 H 2 , are produced. 

Under certain circumstances carbon can also unite directly with 
nitrogen when a substance— e.g. potassium—is present with which 
the compound of the two elements (cyanogen) can combine. From 
carbon, nitrogen, and potassium carbonate we can thus obtain 
potassium cyanide. 

The diamond is always found in alluvial deposits, which have 
been produced by the disintegration of older rock masses, and from 
which it is obtained by a process of washing. Diamond-fields , or 
districts which yield diamonds, occur especially in Brazil, India, 
and at the Cape, as well as in California, Borneo, and in the Ural 
Mountains. Diamonds have also been found in Brazil in a matrix 
of a variety of talc schist rich in quartz—termed itacolumite , and 
distinguished by the ease with which thin strips of it may be bent. 

. But whether this is the original matrix of the diamond, or whether 
the itacolumite was built up from older rock with the diamonds, 
remains as yet undecided. 

The diamond crystallizes in the regular system, and its faces 
are often curved. The crystals may be readily fractured parallel 
to the faces of the octahedron ; it is distinguished by its great 
hardness and bright lustre. Diamonds are usually transparent 
and colourless, or faintly yellow. They are sometimes found of a 
bright yellow or brown colour, and sometimes even blue, green, 
and black. 

Diamonds used as jewels always require cutting and polishing, 
which, on account of their great hardness, can only be done with 
diamond dust. Small diamonds useless for other purposes— 


280 


Text-Book of Inorganic Chemistry. 


diamond-boart —are crushed in a steel mortar. Notwithstanding 
the great hardness of the diamond, it is tolerably brittle, and can 
therefore be easily reduced to powder. 

The form which the diamond receives when ground for a gem 
depends partly upon its original form, and partly upon the use to 
which it is to be put. The most valuable form is that known as the 
brilliant (fig. 52), which may be described as two cones placed base 
to base, of which the upper one is cut off at about half its heigh't. 
Both pyramids are made up of a number of triangular or oblong 



Fig. 52. 


Fig. 53 - 


faces (facets). The brilliant form is always held by its centre 
parts in the setting, so that both upper and lower pyramids are 
free. Differing from this is the table or rosette form (fig. 53), 
consisting of a flat base, on which rises a circle of oblong facets, 
supporting a hexagonal pyramid. 

The value of a diamond depends partly on its size, partly on its 
transparency (water), and partly on its play of colours. The weight 
of a diamond is expressed in carats (0-205 gramme, or 3-165 grains), 
but the value of a diamond, other things being equal, is not pro¬ 
portional to its weight, but to the square of its weight. Thus, if 
one diamond weighs ten times as much as another, its value will 
be approximately a hundred times as much. 

Large diamonds of 100 to 300 carats (20-5 to 61-5 grammes) are 
very rare, and are mostly crown jewels. To these belong the so- 
called Star of the South, found in Brazil, and weighing 127 carats, 
and the Koh-i-Noor, of 106 carats, found in India, and one of the 
English crown jewels. Of the two largest diamonds known, one 
weighs 194 carats, is of a yellowish colour, and is fixed at the end 
of the Russian sceptre; the other, weighing 277 carats, is in the 
possession of the Nizam of Hyderabad. 

Of the formation of the diamond nothing for certain is known, 
except that it is not produced at a high temperature. Its optical 





Carbon. 


281 


properties and microscopical observations seem to show that the 
diamond is of organic origin. All attempts to prepare diamonds 
artificially have as yet given no certain result, although it is always 
more probable that a method might be discovered for converting 
amorphous carbon or graphite into diamond, or of preparing the 
diamond in some other way, than that the problem of the alche¬ 
mists—the transmutation of common metals into gold—should 
ever be solved. 

The present high price of diamonds is due to their scarcity. 
Should larger quantities of diamonds ever be discovered, or should 
it ever be possible to prepare them artificially, their value would be 
considerably diminished. 

On account of its hardness, the diamond is largely used for 
cutting hard stones, and for scratching and cutting glass. For 
simply writing on glass, a splinter of diamond is sufficient, but if 
the glass is to be cut, a deeper scratch is required, and a natural 
convex edge of the diamond is necessary. 

The diamond requires a higher temperature for its combus¬ 
tion than amorphous carbon, but burns easily in the oxy-hydrogen 
blowpipe with excess of oxygen, or when raised to redness by a 
spiral of platinum wire connected with a galvanic battery in an 
atmosphere of oxygen. In the finely divided state the diamond 
bums readily in the air when heated on a piece of platinum foil 
before the blowpipe. In absence of oxygen, for example, in an 
atmosphere of hydrogen, the diamond may be heated to a high 
temperature without any appreciable change. 

Graphite , plumbago , or blacklead , occurs in nature in nodular 
or columnar masses, and occasionally in small tablets belonging to 
the hexagonal system. It is of a dark grey colour, opaque, a good 
conductor of electricity, feels greasy to the touch, and is so soft 
that it easily marks paper. Its specific gravity varies from 2*5 to 
i*8, according as it is more or less pure. Graphite usually occurs 
in isolated spots, but often in considerable quantities. The mines 
of Cumberland and Westmoreland are the oldest, but are now 
nearly exhausted. It is also found in Passau and other parts of 
Germany and Austria, in Bohemia, Greenland, Sicily, and Ceylon, 
and, in large quantities of great purity, in California. Graphite 
may be artificially prepared by solution of carbon in molten iron, 
when it separates out on cooling. 

Finely powdered and pure graphite can be easily compressed 




282 


Text-Book of Inorganic Chemistry. 

at a high temperature to a compact mass of the same specific 
gravity as that occurring in nature. It is only necessary to remove 
air as much as possible by an air-pump before compressing. Thin 
plates or threads for black-lead pencils can be easily sawn from the 
blocks obtained by this process. 

Compact graphite burns almost as difficultly and sometimes , 
more difficultly than the diamond. 

If finely powdered graphite is mixed with three times its weight 
of powdered potassium chlorate, and with enough concentrated 
nitric acid to make a thin paste, then heated for several days to 
6 o°, or placed in direct sunlight until yellow vapours are no longer 
given off, it becomes changed to a bright yellow, transparent, 
crystalline substance of the composition : C n H 4 0 5 . By repeating 
the process several times the whole may be changed into this sub¬ 
stance. It is slightly soluble in water and possesses a faint acid 
reaction, dissolves readily in aqueous alkalies, and in general has 
all the properties of an acid. This peculiar acid is called gra¬ 
phitic acid. 

The graphite which occurs in nature is never pure, but contains 
smaller or larger quantities of inorganic constituents, which remain 
behind as ash on burning. 

Graphite is of considerable value in the arts. It is used for the 
manufacture of black-lead pencils, and of crucibles which are to 
withstand a high temperature, and which are made of a mixture of 
powdered graphite and clay. It is further employed for polishing 
the grains of gunpowder, for coating moulds of gutta-percha or 
plaster to make them conductors of electricity in the process of 
electrotyping, and finally for blacking iron grates and other articles 
of the same metal, which by receiving a coating of the material 
are protected to some extent from rusting. 

A morphous carbon is produced from organic substances when 
they are heated to redness, or when they decay out of contact with 
the air. Wood charcoal may be chosen as representative of 
amorphous carbon which has been produced by the first-named 
process. As its name implies, this substance is obtained by char¬ 
ring wood out of contact with the air, and is prepared either by 
piling logs of wood together, covering with sods, and then slowly 
burning them, or else by heating wood or sawdust in iron retorts. 
By the first process only charcoal is obtained ; but by the second 
and more modern, various volatile products of considerable value, 





Carbon . 


283 

such as acetic acid (pyroligneous acid), wood naphtha, creosote, &c., 
are obtained. According to the kind of wood used the charcoal 
may be more or less porous, and therefore of apparently different 
specific gravity. 

The different varieties of wood charcoal are employed for 
different purposes ; large quantities are used for the manufacture 
of gunpowder. Charcoal is also used (largely in some countries) 
as a fuel. When heated to redness in the air, it bums without 
smoke or flame, and produces a considerably higher temperature 
than the same weight of wood. Wood charcoal is extremely 
porous, and possesses the property of absorbing considerable 
quantities of atmospheric air. Gases which are more easily lique¬ 
fied than common air—for example, ammonia, sulphuretted hydro¬ 
gen, sulphurous anhydride, and carbonic acid—are absorbed by 
charcoal in even greater quantities. It thus possesses the power 
of absorbing noxious gases from the air, and purifies or disinfects 
the air. 

Organic substances which may be suspended or dissolved in 
drinking water and impart a disagreeable taste or smell are removed 
when such water is shaken up with charcoal, or when the water is 
passed through a filter of powdered charcoal. The water which 
runs through such a filter is usually clear, odourless, and pleasant 
to the taste, and charcoal filters are therefore often employed to 
purify drinking water. 

Colouring matter of various kinds is easily removed from a 
liquid by amorphous charcoal—claret, for example, is easily de¬ 
colorized. This property is possessed in a much higher degree by 
the charcoal obtained from animal substances—called animal char¬ 
coal—than by wood charcoal. 

Animal ckarcoal —known also as bone-black, animal-black, 
ivory-black, &c.—is obtained by glowing blood, bones, and other 
animal refuse, excluded from the air. It contains all the inorganic 
substances—principally calcium phosphate and carbonate—present 
in the substances from which it was prepared. These two com¬ 
pounds can be readily extracted by warming the charcoal with 
hydrochloric acid, when, after washing with water, a form of amor¬ 
phous carbon remains, which far exceeds wood charcoal in porosity 
and absorption power. Animal charcoal purified in this way is 
largely used for many technical purposes, especially to decolorize 
the syrup in sugar refineries. 

Wood charcoal contains less inorganic impurity than animal 




284 Text-Book of Inorganic Chemistry. 

charcoal, but is still far from being pure carbon. When burnt 
in the air it leaves behind a considerable ash, rich in potas¬ 
sium carbonate. Pure amorphous carbon may be obtained by 
glowing certain organic substances free from nitrogen and inorganic 
impurities. Pure crystallized sugar or tartaric acid, for example, 
when glowed in a platinum basin and afterwards heated strongly 
in a covered platinum crucible, yields a very pure form of amor¬ 
phous carbon. 

Impure amorphous carbon is found in nature as the various 
forms of coal which have been produced by the decay of vegetable 
matter out of contact with the air, and sometimes also by elevation 
in temperature and strong pressure. All varieties contain, besides 
inorganic constituents—their ash which they leave behind on burn¬ 
ing—varying quantities of hydrogen, oxygen, nitrogen, and sulphur, 
which are mostly driven off when the coal is heated in closed 
vessels. The heavy form of carbon produced by charring coal in 
closed vessels is called coke. Coke is an excellent fuel, and pro¬ 
duces, when freely supplied with air, a higher temperature than 
ordinary coal. It is very largely used for smelting iron and other 
ores. Coke is of a dark grey colour, lustrous, and a conductor of 
electricity. 

The different varieties of coal may be classified according as the 
charring process which they have undergone has been more or less 
complete— i.e. according to their relative proportions of carbon, 
hydrogen, and of oxygen and nitrogen. Coal is usually black in 
colour, mostly of slaty fracture and fatty lustre, and when heated 
in the air generally evolves combustible gases and therefore burns 
with a flame. 

Anthracite is the most highly carbonized form of coal. It is 
distinguished by its hardness and conchoidal fracture ; it pro¬ 
duces very little gas when heated, and therefore burns almost 
without flame. Anthracite is found in considerable quantities in 
South Wales and in Pennsylvania, and is highly valued as a smoke¬ 
less steam-coal and for metallurgical purposes. It does not soil 
the fingers like common coal. 

Bituminous coal in its varied forms comes next to anthracite in 
percentage of carbon. All the commoner kinds of coal belong to 
this class. It contains varying proportions of carbon and hydro¬ 
gen, but always more hydrogen and less carbon than anthracite. 
Bituminous coals rich in hydrogen evolve considerable quantities 
of gas when heated, and are therefore used for the manufacture of 


Ccirboti. 


285 

coal-gas. A variety rich in volatile matter, and especially suited 
for gas making, is cannel coal. This is much more compact in 
structure than ordinary bituminous coal, and, like anthracite, has a 
conchoidal fracture, but its surface is dull, not bright and lustrous 
like common coal and anthracite. 

Brown coal or lignite is always of more recent formation than 
ordinary coal. It is mostly soft and of a brown colour, and often 
exhibits distinct traces of woody structure not to be found in the 
older coals. It contains considerably more hydrogen, oxygen, and 
nitrogen than ordinary coal and less carbon. 

Jet is a compact, black form of lignite, hard enough to take a 
high polish, and largely used for ornamental articles. 

Peat or turf is produced in marshy localities by the gradual 
decay of grasses, mosses, and other marsh plants. It contains 
large quantities of inorganic constituents, and leaves a large 
amount of ash when burnt. It is of much less value as a fuel than 
any of the varieties of coal mentioned above. 

The following table 1 gives the approximate percentage compo¬ 
sition of the organic parts of wood, peat, and the different kinds of 
coal, omitting the ash in all cases :— 



Carbon 

Hydrogen 

Oxygen, 
Nitrogen, 
and Sulphur 

Oak. 

50*17 

6 'c 8 

4375 

Peat. 

Lignite (Bovey Tracey) . 

59"83 

67 *90 

5*78 

5*73 

34*39 

26 *37 

Staffordshire coal . . . ! 

79*37 

5*37 

15*26 

Wigan cannel ccal . . . 1 

82-47 

5' 6 7 

11 •86 

Newcastle caking coal 

87*95 

5*34 

6*71 

South Wales anthracite . 

9* *85 

3*35 

4’8o 


Amorphous carbon in a finely divided state is obtained as soot 
I or lampblack , by the incomplete combustion of volatile organic 
I compounds, rich in carbon, such as turpentine, benzol, or wood 
containing turpentine and resin. If a cold body is placed in the 
flame of a candle or oil-lamp, it soon becomes covered with soot 
or finely divided carbon, just as arsenic is deposited from the flame 
| of burning arsine. Lampblack is used in a more or less pure state 
I for various purposes—the finer varieties for preparing Indian ink, 

! the more coarse as a black paint and for printer’s ink. Unless 
purified by special processes, it always contains varying quantities 
of hydrocarbons mixed with it. 

1 From Watts’s Dictionary of Chemistry. —Ed. 















286 


Text-Book of Inorganic Chemistry. 


COMPOUNDS OF CARBON. 

No element exists in so large a number of compounds as carbon. 
By far the largest proportion of these belongs to organic chemistry, 
and we should be justified in referring all compounds of carbon to 
this division, if it were possible to separate inorganic and organic 
compounds distinctly from one another, and if it were not abso¬ 
lutely indispensable to describe some of the more simple of the 
carbon compounds under inorganic chemistry. 

To these belong carbonic acid and oxide, oxalic acid, carbon 
disulphide and oxysulphide, carbon tetrachloride, marsh gas and 
cyanogen, together with hydrocyanic acid, the cyanides, and a few 
other compounds. 


CARBONIC ANHYDRIDE, or CARBONIC ACID. 1 

Composition : C0 2 . 

Carbonic acid, which is always produced when carbon in any 
form or any compound of carbon is burnt with a free supply of air, 
is a colourless gas, of peculiar prickly smell and faintly acid taste. 

It occurs very widely distributed in nature, both free and com¬ 
bined with various bases. It is a normal constituent of atmo¬ 
spheric air (p. 200), and is contained in most natural waters in the 
free state. In immense quantities combined with lime it forms the 
different natural varieties of calcium carbonate—common limestone , 
calcspar , arragonite , marble , chalk , &c. ; further, as magnesium car¬ 
bonate in magnesite , as double carbonate of calcium and magnesium 
in dolomite , as ferrous carbonate in spathic iron ore , as barium and 
strontium carbonates in witherite and strontianite respectively, 
and, finally, as sodium carbonate in trona or soda— to mention 
only the more important compounds. 

In many places, carbonic acid is evolved in considerable quan¬ 
tities from clefts in the earth, usually in the neighbourhood of 
active or extinct volcanoes. The gas is, of course, invisible, but 
can be easily recognized by chemical means. Among the most 

1 True carbonic acid, which would result from the union of the anhydride 
with water, is unknown, and the anhydride (C0 2 ) is therefore usually called 
carbonic acid, for the sake of brevity.—E d. 



Carbonic Acid. 


287 

celebrated localities of these carbonic acid springs are the Grotto 
del Cane, near Naples, those in Pyrmont, and in various parts of 
Germany, and the well-known poison valley of Java. Generally, 
the carbonic acid carries water with it in its upward course to the 
surface of the earth, and often with considerable force, owing to the 
pressure of the gas. At Nauheim a spring rises 80 feet high, probably 
owing to the force with which the compressed carbonic acid escapes. 

Finally, carbonic acid is produced in such large quantities 
during the fermentation of sugar in the manufacture of alcoholic 
liquors, that attempts have been made to utilize the gas—for ex¬ 
ample, in the manufacture of white-lead. 

It is thus clear that there is no lack of material from which 
carbonic acid may be prepared; but if the gas is wanted pure, the 
choice of substances becomes more limited. All limestones which 
contain traces of organic matter when treated with strong acids 
yield carbonic acid with an unpleasant odour and taste, because 
small quantities of these organic substances are carried away with 
the gas. And carbonic acid prepared from such material should 
therefore never be used for the manufacture of mineral waters. On 
the other hand, the gas evolved from calc-spar, marble, magnesite, 
and many kinds of dolomite and limestone, as well as from acid sodium 
carbonate, is always pure and free from any foreign smell or taste. 

For the preparation of small quantities of carbonic acid Kipps 
apparatus is best adapted, especially as with this arrangement a 
stream of pure carbonic acid can be obtained at any time by simply 
opening a stop-cock. Carbonic acid may be evolved from mag¬ 
nesite or acid sodium carbonate by means of dilute sulphuric acid, 
and need then only be dried (when necessary) by passing through 
concentrated sulphuric acid. To prepare carbonic acid from dolo¬ 
mite or calcium carbonate, it is better to use hydrochloric acid, as 
sulphuric acid produces difficultly soluble calcium sulphate, which 
surrounds the pieces of marble or dolomite and tends to prevent 
their contact with the acid. If sulphuric acid is preferred, the 
limestone or dolomite must be finely powdered, and the mixture 
continuously agitated. The process is represented by the following 
equation :— 

CO * 0 2 Ca + S 0 2 ( 0 H) 2 = S 0 2 * 0 2 Ca + H 2 0 + C 0 2 . 

When carbonic acid is liberated from limestone or marble with 
dilute hydrochloric acid, calcium chloride then remains behind in 
solution, and the equation becomes :— 

C 0 - 0 2 Ca + 2HCI = CaCl 2 + H 2 0 + C 0 2 . 


288 


Text-Book of Inorganic Chemistry. 


The gas prepared in this way carries some of the volatile acid 
with it, and should be freed from this by passing through several 
wash-bottles filled with water. 

Carbonic acid may be collected over mercury or over water, in 
which it is only slightly soluble. At the ordinary temperature 
water dissolves about its own volume of the gas. Since, however, 
carbonic acid is considerably heavier than air, vessels are best filled 
with the gas, especially when required dry, by leading the gas with 
a long glass tube to the bottom of the vessel to be filled (fig. 54). 
The carbonic acid then gradually displaces the lighter air until in 
a short time the whole vessel is filled with the gas. 



Fig. 54 - 


Carbonic acid has a specific gravity of 1-52, corresponding to a 
molecular weight of 1*52 x 28-88 = 44, and its molecule is therefore 
C 0 2 . It is thus considerably heavier than air, and diffuses from 
an open jar so slowly that it can be poured like water from one 
vessel to another. If two jars are taken, one filled with carbonic 
acid and the other with air, the carbonic acid can, with care, be 
poured almost completely from one into the other, as may be easily 
shown by introducing a burning taper. 

Carbonic acid neither burns nor supports combustion. A 
burning candle is at once extinguished in a jar of the gas; even if 




















Carbonic Acid. 289 

the air only contains a small quantity of carbonic acid, a candle 
ceases to burn. Respiration in such an atmosphere is also impos¬ 
sible, and carbonic acid, without being actually poisonous, is fatal 
to the life of animals, because it takes the place of the oxygen 
necessary to aerate the blood. Death therefore ensues simply from 
suffocation. That carbonic acid is not a true poison is at once 
evident from the following considerations : we continuously inhale 
carbonic acid, largely diluted it is true, into our lungs ; carbonic 
acid is continually produced in considerable quantities by the slow 
combustion going on in our bodies, and forms a large proportion 
of the air exhaled from our lungs ; finally, continued use of effer¬ 
vescing drinks, rich in carbonic acid, does not apparently produce 
any injurious effect. 

Nitrogen, like carbonic acid, neither burns nor supports com¬ 
bustion, nor can it be respired ; the latter gas may, however, be 
distinguished from the former by the fact that when brought into 
contact with clear lime-water it at once produces a white tur¬ 
bidity or precipitate of calcium carbonate, insoluble in water, and 
which again dissolves with effervescence in dilute hydrochloric 
acid. Baryta-water is even better adapted than lime-water for the 
detection of carbonic acid. 

Carbonic acid possesses the properties of a weak acid anhy¬ 
dride, and unites therefore with bases. The combination, however, 
takes place slowly, and but little heat is evolved. If a glass tube 
closed at one end is filled with carbonic acid and the open end 
placed in a vessel of caustic potash, the liquid rises in the tube 
slowly, more quickly if shaken, until ultimately the whole of the 
carbonic acid disappears. We employ this property of the gas not 
only to separate it when mixed with other indifferent gases, but 
also to determine its amount quantitatively. 

The quantity of carbon present in an organic substance is 
found by heating it with some substance that readily gives up 
oxygen {e.g. copper oxide) in a tube of hard glass. By this process 
the hydrogen present is converted into water, and the carbon into 
carbonic acid. These compounds are allowed to pass first through a 
U-tube containing calcium chloride to absorb the water, and then 
through a series of bulbs (called Liebig’s potash-bulbs, fig. 55), 
which contain strong caustic potash, and absorb the carbonic acid. 
If the calcium chloride tube and the potash bulbs are separately 
weighed before and after the experiment, the increase in weight of 
the former gives the weight of water formed, and of the latter the 

U 


290 Text-Book of Inorganic Chemistry. 

weight of carbonic acid. From these weights the actual weights 
of hydrogen and carbon present in the substance can be easily 
calculated. 

The affinity of carbon for oxygen increases as the temperature 
rises, and at a white heat carbon can 
even extract oxygen from potassium 
oxide or carbonate and set free metallic 
potassium. But at a red-heat, and when 
the carbonic acid is in excess, potassium 
can abstract oxygen from carbonic acid. 
If a small piece of potassium is heated 
in a bulb-tube of hard glass to low red¬ 
ness while a stream of dry carbonic acid 
is passed over it, the fused metal soon 
begins to burn with a dark red light, 
and becomes covered with a coating of 
potassium carbonate and amorphous 
carbon. The following equation explains the process :— 



4K + 3C0 2 = 2CO(OK) 3 + C. 

This experiment shows that colourless carbonic acid contains 
black charcoal, and that it can be reduced to this substance. 

Carbonic acid is the anhydride of an unknown dibasic acid, 


the composition of which would be 



but which has not 


yet been prepared. The majority of its salts are insoluble in water ; 
the only soluble carbonates being those of the alkali metals : 
potassium, sodium, lithium (ammonium, rubidium, and caesium), and 
of thallium. Normal potassium carbonate (potashes) and normal 
sodium carbonate (soda) possess an alkaline reaction, because 
the strong basic properties of these oxides are not completely 
neutralised by the weak carbonic acid. Its acid salts of these 


(OH 


two metals— eg. acid sodium carbonate : —react neutral. 

When carbonic acid is led into aqueous ammonia the chief pro¬ 


duct is ammonium carbonate, of the composition : CO 


onh 4 . 

ONH 4 *~“ 


C 0 2 + 2 NH 3 + H 2 0 = CO pH, 

This salt cannot, however, be produced from dry ammonia 
and dry carbonic acid, because the molecule of water necessary 



Carbonic Acid. 


291 


for its production is now absent. The two gases unite, however, 
with considerable evolution of heat, and produce a salt, which 

[ NH 0 


[onh 4 


9 


is ammonium carbamate, having the composition: CC)j| 

(p. 186). 

Carbonic acid can be condensed by a high pressure to a liquid 
which, when allowed to evaporate in the air, produces a sufficient 
diminution in temperature to freeze it. For the 
preparation of liquid carbonic acid, the apparatus ^ 

devised by Natterer is now almost universally em¬ 
ployed. This consists of a compression pump, 
which also acts as a suction pump, and a strong 
wrought-iron vessel. The latter vessel (fig. 56) is 
screwed at its base to the top of the cylinder of the 
pump, and is there provided with a small valve 
opening upwards at the end of every stroke of the 
piston, and otherwise kept closed by a weak spiral 
spring. This valve, therefore, only allows gas to enter 
the strong vessel, and prevents any from leaving it. 

At the upper end .of this vessel, a narrow tube, 
is screwed on at right angles, which can be set in 
communication with the interior of the vessel by 
raising a conical valve attached to the screw, t. 

This tube is to allow the liquid carbonic acid to 
escape after about 250 grammes have been pumped 
in. For this purpose, the receiver is unscrewed, 
turned upside down, and the screw, /, carefully 
turned. 

If a jet of liquid carbonic acid is allowed to 
flow out of the side tube into a metallic vessel, 
which is so arranged that it receives a circular 
motion, and therefore rapidly evaporates, a large 
portion of the carbonic acid freezes at the low tem¬ 
perature produced, and fills the vessel as a light, snow-like mass. 

Condensed carbonic acid is a colourless, transparent, mobile 
liquid of 0-95 specific gravity at o°, and only 078 at + 25° It 
therefore expands under the influence of heat not only greater 
than any other liquid, but also greater than the gases—a very 
remarkable property. 

The tension of liquid carbonic acid at different temperatures 
is about as follows :— 


Fig. 56. 








292 


Text-Book of Inorganic Chemistry. 


At - 79 ° 


)> 

5 ) 


)5 


)> 

» 


~ 3 o° 
— io° 
o° 

+ IO° 

+ 30 0 
+ 35 ° 


i-i atmospheres. 
i8 „ 

27 » 

36 „ 

46 » 

74 >» 

80 „ 


During free evaporation in the air its temperature falls to about 
— 79 0 . 

The solid, snow-like carbonic acid evaporates but slowly even 
on a warm hand, and notwithstanding its low temperature produces 
scarcely a feeling of cold. It does not, in fact, touch the hand at 
all, but is separated by a layer of the gas which it is continually 
giving off. The phenomenon is similar to that produced when a 
drop of cold water is allowed to fall into a red-hot platinum dish ; 
the drop then remains suspended by its own vapour and does not 
touch the hot metal at all. If, however, a piece of solid carbonic 
acid is pressed firmly on the skin, a burning pain is felt and a 
blister is raised on the spot touched. The action is, in fact, the 
same as when the skin is burnt with a hot body. 

If solid carbonic acid is mixed with some volatile liquid—^, 
ether—both evaporate with great rapidity, the mixture begins to 
boil, and a sufficiently low temperature (about - ioo c ) is produced 
to freeze mercury in a few seconds. 

Carbonic acid is largely used for the manufacture of artificial 
mineral waters. Pure water in which various salts are dissolved, 
according to the mineral water required, is impregnated with 
carbonic acid in thick metallic vessels. The gas is pumped in by 
a special apparatus, and the liquid is kept constantly agitated. In 
this way not only can many mineral waters be imitated, but other 
compositions of any required kind can be prepared. Water in 
which small quantities of salts have been dissolved retains its 
carbonic acid when exposed to the air longer than pure water. 

Carbonic acid is also used in the arts for the preparation of 
white-lead, and of salicylic acid from phenol. Liquid carbonic 
acid is now employed on a large scale by Krupp at Essen, in Ger¬ 
many, for compressing cast-steel during solidification. 







Carbamic Acid. 


293 


CARBAMIC ACID. 
Composition of the hypothetical acid : CO 


nh 2 

OH 


As previously mentioned (p. 290), dry carbonic acid and dry 
ammonia unite to form a salt which cannot be ammonium car¬ 
bonate, as the molecule of water necessary for the formation is 
wanting. 

The compound produced is the ammonium salt of a monobasic 
acid, which is as little known in the free state as the true carbonic 


acid : 


CO 


j OH 
) OH' 


It has the same relation to this dibasic acid as 


chlorsulphonic acid (p. 165): S0 2 j has to sulphuric acid, 

except that carbamic acid contains an atom of amidogen, instead 
of one of chlorine, in place of an atom of hydroxyl. 

The production of ammonium carbamate may be easily under¬ 
stood if we imagine one of the two molecules of ammonia which 
unite with one molecule of carbonic acid to give up an atom of 
hydrogen to the other. By this means amidogen is produced on 
the one hand, and ammonium on the other, which then unite with 
the carbonic acid as shown by the following equation :— 


C0 2 + 2 NH 3 


(NH 

|onh 4 


Ammonium carbamate is easily obtained by leading carbonc 
acid and ammonia, both perfectly dry, into absolute alcohol, 
when the salt separates as a white, voluminous, crystalline mass. 
When heated with absolute alcohol in sealed tubes to ioo° 
and allowed to cool slowly, the salt separates out in large thin 
crystalline plates. 

This salt is always contained together with ammonium car¬ 
bonate in the commercial carbonate of ammonia. In contact with 
water it gradually takes up a molecule of this substance, and after 
some time is converted into ammonium carbonate. It is distin¬ 
guished from this latter compound by the fact that its aqueous 
solution when mixed with a solution of calcium chloride does not 
at once deposit a white precipitate of calcium carbonate, because 
the calcium carbamate which is first produced is soluble in water. 


Calcium carbamate : (C 0 NH 2 ) 2 0 2 Ca + H 2 0 , is produced when 


294 Text-Book of Inorganic Chemistry. 

carbonic acid is led into a mixture of thick milk of lime and four 
times its volume of strong ammonia. The filtered liquid is then 
cooled to o° and precipitated with absolute alcohol : the amorphous 
precipitate thus produced soon changes into small prismatic crystals 
of the above composition. Calcium carbamate is tolerably soluble 
in water, but the solution soon becomes turbid from the calcium 
carbonate produced by its decomposition with water. In aqueous 
ammonia the salt remains longer undecomposed. When gently 
heated it loses its water of crystallization, and then bears a tempe¬ 
rature of i8o° without decomposition. 

Carbamic acid forms more stable organic compounds with the 
alcoholic radicals, producing then a series of bodies called urethanes.^ 

Ammonium carbamate, which contains the constituents of one 
molecule of water less than ammonium carbonate, contains itself 
one molecule of water more than carbamide, thus :■ 


Ammonium carbonate . 

. . CO 

(ONH 4 

|onh 4 

Ammonium carbamate . 

. CO 

(ONH 4 

(NH 2 

Carbamide or urea 

. CO 

(NH 3 

NH„ 


Carbamide not yet prepared in the free state is isomeric with 
urea, a normal constituent of urine ; it may be considered either as 
the hypothetical carbonic acid in which both atoms of hydroxyl 
have been displaced by amidogen, or as carbonic anhydride in 
which one atom of oxygen is displaced by two of amidogen. Am¬ 
monium carbamate is converted with loss of water into carbamide, 
which at once changes into urea, when heated in sealed tubes up 
to 140°. 


CARBONIC OXIDE. 

Composition : CO. 

The second lower oxide of carbon, which contains the element 
as a dyad, is, like carbonic acid, a colourless gas, but is distinguished 
from this compound principally by its combustibility. When heated 
in the air it catches fire and burns with a blue flame, forming car¬ 
bonic acid. 

Carbonic oxide has a specific gravity of 0*968, corresponding to 



Carbonic Oxide. 


295 


a molecular weight of 28, is scarcely soluble in water, and is only 
condensed to a liquid under the greatest pressure combined with 
a very low temperature. The gas, although tasteless and odour¬ 
less, is very poisonous ; air containing only a few per cent, of it 
when breathed produces giddiness and intense headache, followed 
by insensibility, and after a little time by death. 

Carbonic oxide does not occur free in nature like carbonic acid, 
but is always produced by the incomplete combustion of carbon or 
carbonaceous substances— i.e. when these are burnt with an insuffi¬ 
cient supply of air. It may also be prepared from carbonic acid if 
this gas is led over glowing charcoal. The blue flames seen when 
air flows freely over large masses of glowing carbon—for example, 
when the door of the firebox of a steam-engine is opened—are due. 
to burning carbonic oxide, produced by the contact of carbonic acid 
with the red-hot carbon. 

For the preparation of carbonic oxide it is best to employ either 
oxalic or formic acid, or else a salt of the last-named acid. 

Oxalic acid, which occurs in nature in many plants, especially 
the sorrels, and which can also easily be prepared artificially, is 
a solid crystalline substance of strong acid properties, and having 


the composition j 


CO-OH 

CO-OH’ 


This body, when mixed with concen 


trated sulphuric acid and warmed, breaks up into water, carbonic 
acid, and carbonic oxide :— 


(CO-OH 

{CO-OH 


= H 2 0 


+ co 2 + CO. 


The sulphuric acid first removes the water of crystallization from 
the oxalic acid, and then, on account of its strong attraction for 
water, abstracts a further two atoms of hydrogen and one of oxygen, 
which unite to form water. As, however, the oxalic anhydride : 

CO f O) w ^ich is thus produced, does not appear to exist, this hypo¬ 
thetical substance at once breaks up into carbonic acid and car¬ 
bonic oxide :— 

£0)o = C0 2 + CO. 

In order to purify the carbonic oxide from the large quantities 
of carbonic acid mixed with it, the mixed gases must be passed 
through a system of wash-bottles containing strong caustic potash, 
which absorbs the carbonic acid, but allows the carbonic oxide to 


296 Text-Book of Inorganic Chemistry . 


pass on unchanged. If the gases are evolved rapidly, it is difficult 
to separate the whole of the carbonic acid from the mixture. It is, 
therefore, better to prepare the carbonic oxide from formic acid, 
which breaks up when heated with sulphuric acid into carbonic 
oxide and water only. 

(H 

Formic acid is a monobasic acid of the composition : j CO-OH 


and its sodium salt 


• J H 

' (CO-ONa* 


On heating this compound 


with concentrated sulphuric acid, formic acid is first set free, and 
this is then decomposed into carbonic oxide and water. The fol¬ 
lowing equation expresses the entire reaction :— 


( H (OH 

(CO-ONa + (OH 


S ° a {oNa. + C0 + H *°- 


The water produced is, as before, retained by the excess of sulphuric 
acid. 

This process easily yields considerable quantities of pure car¬ 
bonic oxide, and is the best method to employ, now that formic 
acid is artificially prepared in large quantities. 

Carbonic oxide is an indifferent substance in so far that it 
neither unites with acids nor with bases to form salts. And not¬ 
withstanding the tendency of the unsaturated dyad carbon which 
it contains to pass into saturated tetrad carbon by union with some 
other element, it neither unites with oxygen, nor with sulphur, or 
chlorine at the ordinary temperature. Combination with oxygen 
to form carbonic acid or with sulphur to form carbon oxysulphide : 
COS, requires a high temperature ; and to produce its compound 
with chlorine—carbon oxychloride : COCl 2 —either a high tempe¬ 
rature or direct sunlight is necessary. 

Although carbonic oxide does not unite with aqueous caustic 
potash, the solid substance combines easily and completely with 
the gas when the two bodies are heated up to ioo° or over. The 
sole product of this union is potassic formate :— 

KOH + CO = |» 0 . 0K 

We can, therefore, obtain carbonic oxide from formic acid by the 
abstraction of a molecule of water, and reconvert the gas into 
formic acid by the action of caustic potash. 

If carbonic oxide is led into a concentrated hydrochloric acid 
solution of cuprous chloride : Cu 2 Cl 2 , it is absorbed in considerable 


Carbon Oxychloride. 297 

quantities, and a crystalline compound is produced which contains 
the elements of cuprous chloride, carbonic oxide, and water. 


CARBON OXYCHLORIDE (Carbonyl Chloride). 

Composition : COCl 2 . 

This compound, also known as phosgene gas, because formed 
by the action of light, may be considered as carbonic acid in which 
one atom of oxygen has been displaced by two atoms of chlorine, 
and has the same relation to this substance as sulphuryl chloride 
(p. 164) to sulphuric anhydride. It is produced, without explosion, 
when equal volumes of carbonic oxide and chlorine are exposed to 
direct sunlight, the volume of the mixed gases diminishing to one 
half. 

Carbon oxychloride is a colourless gas of peculiar, suffocating 
odour, easily condensed in an ordinary freezing mixture to a 
colourless liquid, with a specific gravity of 1-43, and boiling at + 8°. 
The specific gravity of the gas is 3*46, and its molecular weight 
3-46 x 28*88 = 99*9, corresponding to the formula : COCl 2 . Its com¬ 
position is also proved from the fact that two volumes (one 
molecule) of carbonic oxide and two volumes of chlorine produce 
two volumes of the gas :— 

CO + Cl 2 = COCl 2 

2 vols. 2 vols. 2 vols. 

In contact with water it is at once decomposed into carbonic acid 
and hydrochloric acid. 


OXALIC ACID. 

„ . ... (CO-OH 

Composition . j (^q . 014 

Oxalic acid occurs in nature as the acid potassium salt in the 
various kinds of sorrel and other plants, and also in combination 
with lime as calcium oxalate. 

It crystallizes in colourless rhombic prisms with two molecules 
of water of crystallization, which can be easily expelled on heating. 
Oxalic acid possesses a strong acid reaction, and is tolerably 




298 Text-Book of Inorganic Chemistry . 


soluble both in water and alcohol. Nine parts of water dissolve 
one of the acid at the ordinary temperature, and a much larger 
quantity when hot. When the anhydrous acid is heated above 
ioo°, a portion sublimes unchanged, but the greater part is decom¬ 
posed into carbonic acid and formic acid 


(CO-OH 
(CO-OH “ 


+ 


(H 

(CO-OH 


which latter then further decomposes into water and carbonic 
oxide. 

Oxalic acid is not now obtained, as formerly, from the acid 
juice of the sorrel, nor by the expensive process of oxidizing cane- 
sugar with nitric acid, but is exclusively prepared from some form 
of cellulose, generally sawdust. The sawdust is mixed with caustic 
potash, to which a certain, not too large a quantity of caustic 
soda may be added, and the mass heated until it fuses. The 
cellulose is thus destroyed—oxidized—with evolution of hydrogen, 
and a large part of its carbon converted into oxalic acid, which 
unites with the potassium and sodium. 

The fused mass is then extracted with water, neutralized with 
hydrochloric acid, and mixed with a solution of calcium chloride. 
The insoluble calcium oxalate which is thus produced is filtered 
off, suspended in water, and digested with the requisite quantity of 
sulphuric acid. By this means slightly soluble calcium sulphate 
(gypsum) is formed, and the oxalic acid remaining in solution is 
. deposited on evaporation ; the crude acid being afterwards purified 
by repeated crystallization. 

Of theoretical interest is the fact that oxalic acid can be directly 
obtained from carbonic acid by reduction. If dry carbonic acid is 
led over potassium amalgam, a brisk reaction takes place, and 
potassium oxalate is formed. The mercury of the potassium 
amalgam undergoes no change, but serves simply as a diluent :— 


2C0 2 + 2K 


fcCO-OK 

ICO-OK 


Oxalic acid is a dibasic acid, and is a chemical compound of two 
atoms of the monad radical oxatyT. CO-OH, just as a molecule 
of hydrogen consists of two separate atoms. The oxalates are 
mostly insoluble in water. One of the most insoluble is calcium 
oxalate, which is not even dissolved by acetic acid, and which 
serves, therefore, for the detection and estimation of oxalic acid 
{and calcium). 


Oxalic Acid. 


299 


Oxalic acid in its relations to nitric acid is very stable; even 
-when the two are boiled together it is only slowly oxidized to 
carbonic acid. Other oxidizing agents, as potassium permanganate 
or manganese peroxide, convert it easily and completely into car¬ 
bonic acid and water. 

On heating oxalic acid to the point when it begins to sublime, a 
portion of it is converted into carbonic acid and formic acid, as 
previously explained j and if the oxalic acid is dissolved in aqueous 
glycerine, and then heated to ioo°, the change into these two sub¬ 
stances is so complete and easy that this is employed as the best 
method for preparing formic acid. 

The conversion of oxalic acid into carbonic acid, carbonic oxide, 
and water, when heated with concentrated sulphuric acid, has been 
already referred to (p. 295). 


Oxamic Acid and Oxamide. 


Just as the hypothetical carbonic acid becomes carbamic acid 
when one of its atoms of hydroxyl is displaced by amidogen, and 
carbamide on displacement of both atoms by the same radical, so 
the same change in oxalic acid produces oxamic acid and oxiamide 
respectively :— 

~ (CO-OH 

Oxalic acid.(CO-OH 

^ (CO-NH 2 

Oxamic acid.{CO-OH - 

^ .j j CO-NHo 

Oxamide.\CO-NH, 


Oxamic Acid'. (cO-OH* is a white > crystalline powder, slightly 
soluble in water. It is produced by heating acid ammonium 


oxalate:— 

(CO-ONH 4 
(CO -OH 


(CO-NH 3 
{ CO • OH ~ + 


H 2 0 ; 


or by boiling oxamide with aqueous ammonia 


(CO-NH 2 
(CO-NH 2 + 


H 2 0 


j CO-NH,, 
(CO-ONH 4 


and is separated from the solution of its ammonium salt thus 
obtained by hydrochloric acid. 

Oxamic acid is a monobasic acid, and is easily dissolved by 


300 Text-Book of Inorganic Chemistry. 

alkalies, but when boiled with these it takes up the elements of 
water and is converted into oxalic acid and ammonia. Boiled with 
water, it gives acid ammonium oxalate. 


Oxamide : jco-NH* * s a w ^ te crystalline powder, tasteless 

and odourless, insoluble in cold water or alcohol, and only slightly 
dissolved by these liquids when hot. It is produced when normal 
ammonium oxalate is heated :— 


JCO-ONH 4 
(CO-ONH 4 “ 


(CO *NHn 


2 H 2 0 . 


\co-nh 2 

It is, however, better prepared from an organic compound of 
oxalic acid—oxalic ether, a colourless liquid with pleasant ethe¬ 
real odour and insoluble in water 1 —by shaking it with aqueous 
ammonia. The oily liquid gradually disappears, and a crystalline 
powder of oxamide takes its place, with the simultaneous formation 
of alcohol :— 


jCO-OC 2 H 5 

(CO-OC 2 H 5 

Oxalic ether 


+ 


2NH 3 


j CO • NH 0 
(CO-NH 2 

Oxamide 


2C 2 H 5 'OH 

Alcohol 


When carefully heated, oxamide may be sublimed, but quick 
heating decomposes it into various products, among which cyanogen 
may be recognized by its odour. Boiling with aqueous alkalies 
changes it into oxalic acid and ammonia. 


CARBON DISULPHIDE. 

Composition : CS 2 . 

This compound, the anhydride of a sulpho-acid corresponding 
to carboni: acid, was discovered by Lampadius towards the end of 
the preceding century. It is produced in a precisely similar method 
to carbonic acid—viz. by burning carbon in sulphur gas. A con¬ 
siderably higher temperature is, however, necessary to cause carbon 
to burn in sulphur than in oxygen. 

Carbon disulphide is a colourless, mobile, very volatile liquid, 
insoluble in water, in which it sinks, refracting light strongly, and 

1 Oxalic ether is prepared by distilling a mixture of dry oxalic acid and 
absolute alcohol. 



Carbon Disulphide. 301 

possessing an unpleasant ethereal odour. It has a specific gravity 
of 1*29, boils at 46°, and is miscible with alcohol and ether in all 
proportions. It is very inflammable, and burns with a bluish flame 
to carbonic acid and sulphurous acid, with the separation of sul¬ 
phur if the air supply is insufficient. 

Carbon disulphide is prepared by heating freshly glowed wood- 
charcoal in large tubulated earthenware retorts to bright redness 
in a suitable furnace, and dropping in pieces of sulphur from time 
to time. The neck of each retort is connected air-tight with a large 
receiver in which the uncombined sulphur condenses, and from 
this the carbon disulphide vapours are led through well-cooled 
condensers. The crude product contains sulphur, sulphuretted 
hydrogen, and other impurities, imparting to it a very unpleasant 
odour, and from which the carbon disulphide cannot be completely 
freed even by repeated redistillation. It may, however, be effec¬ 
tually purified by shaking with mercury and with mercuric chloride, 
and then distilling. 

Carbon disulphide is employed for numerous technical purposes, 
and is now manufactured on a large scale and sold at a low price. 
It is an excellent solvent for many substances— e.g. for iodine, sul¬ 
phur, phosphorus, fatty and ethereal oils, resins, &c.—and in con¬ 
sequence of this property is largely used to impregnate india-rubber 
with sulphur in the manufacture of vulcanised rubber, to free the 
finer kinds of wool from fat, and to extract fatty oils from seeds 
containing them. 

Air containing small quantities of the vapour of carbon disul¬ 
phide, when respired, acts in a similar manner to chloroform and 
produces insensibility; in laiger quantities it is poisonous. At the 
same time carbon disulphide possesses strong antiseptic properties ; 
it prevents the putrefaction of meat, and stops the processes of 
fermentation. 

Pure carbon disulphide exposed to sunlight soon becomes of a 
yellow colour and acquires the unpleasant odour of the crude sub¬ 
stance. After some time it deposits a brown amorphous substance, 
which apparently has the composition of carbon monosulphide: 
CS, and which is insoluble in carbon disulphide and most other 
liquids. It is probable that the light decomposes the carbon 
disulphide into sulphur and the still unknown monosulphide, which 
then polymerises to form this brown substance. 

Dry chlorine gas decomposes carbon disulphide with formation 
of disulphur dichloride and carbon tetrachloride. 


302 


Text-Book of Inorganic Chemistry. 


Carbon disulphide is the anhydride of a dibasic sulpho*acid 

sulphocarbonic acid : CS | and as SUch Can Unite with aqueoUS 

solutions of sodium or potassium sulphides, or other sulphides 
soluble in water, to form sulphosalts— e.g. with potassium sulphide 

to form the compound : CS -j g —potassium sulphocarbonate. 

The sulphocarbonates can only be obtained in the solid form 
with difficulty, as they very easily decompose. The potassium 
compound separates out from its concentrated solution in yellow 
deliquescent crystals, containing water. 

Aqueous alkalies also dissolve carbon disulphide and produce 
a mixture of sulphocarbonate and carbonate :— 

3 CS 2 + 6KOH = 2CS {sK + CO {oK + 3H 2 0 * 


ammonia is used, ammonium sulphocarbonate : 

^ t 3 4 is also produced, together with ammonium sulphocar- 
oN ±1^ 


If aqueous 
|SNH 4 ; 


bamate : CS j g^ and ammonium sulphocyanate : CSN • NH 4 . 

Alcoholic potash when warmed with carbon disulphide dissolves 
considerable quantities, and after a short time the liquid almost 
solidifies to a mass of small yellow crystals. This compound is 
not potassium sulphocarbonate, but is the potassium salt of an 
organic acid, termed xanthogenic acid. 

If a sulphocarbonate is mixed with alcohol and then hydro- 

r (SH 

chloric acid added, free sulphocarbonic acid : CS j gH separates 

as a dark yellow oil, with a disagreeable odour. The acid readily 
decomposes into carbonic disulphide and sulphuretted hydrogen. 


CARBON OXYSULPHIDE. 

Composition : COS. 

This compound, which may be considered as carbonic acid in 
which one atom of oxygen is displaced by sulphur, and is therefore 
the intermediate compound between carbonic acid and carbon 
disulphide, is a colourless gas, with an unpleasant odour, combus- 



Carbon Oxysulphide, 303 

tible and easily inflamed. Water dissolves its own volume of the 
gas, but decomposes it on standing into carbonic acid and sulphu¬ 
retted hydrogen. 

Carbon oxysulphide is found in some mineral waters, and may 
he artificially prepared by leading carbonic oxide with excess of 
sulphur vapour through a glass tube heated to low redness, or by 
heating a mixture of sulphuric anhydride and carbon disulphide. 
The best method for its preparation consists in heated potassium 
sulphocyanate with moderately dilute sulphuric acid. There are 
then produced, besides carbon oxysulphide, acid ammonium sul¬ 
phate and acid potassium sulphate as secondary products, together 
with some hydrocyanic acid and sulphuretted hydrogen :— 

CSNK + 2S0 2 )§« + HjO = COS + SO^g^ + SO a |g£ 

Alkalies, like water, decompose carbonic oxysulphide with 
formation of carbonic acid and sulphuretted hydrogen ; thus, with 
potash the change is as follows :— 

COS + 4KOH = CojgK + K 3 S + 2H 2 0. 

A second compound isomeric with carbonic oxysulphide, of the 
composition : CSO, may exist, but has not yet been prepared. 


COMPOUNDS OF CARBON AND HYDROGEN. 

Carbon and hydrogen, although they have little attraction for 
one another in the free state, unite together to form a large series 
of compounds—the so-called hydrocarbons. Nearly all these 
belong to organic chemistry, and we can here only refer to two of 
the simplest—viz. methane and ethylene. 





304 


Text-Book of Inorganic Chemistry. 


METHANE, or MARSH GAS. 

Composition : CH 4 . 

This, the most important compound of carbon and hydiogen, is 
a colourless, odourless gas, insoluble in water, and burning in the 
air with a pale, non-luminous flame. Its specific gravity is 0-55, 
and its molecular weight, therefore, 0-55 x 28*88 = 15-9, correspond¬ 
ing to the above composition. Its low specific gravity gave rise 
tolts old name light carburetted hydrogen, as distinguished from 
heavy carburetted hydrogen , or ethylene. 

Methane is found in nature as a product of the decomposition 
of organic substances out of contact with air—for example, by the 
putrefaction of vegetable matter at the bottom of stagnant pools 
and marshes. The marsh gas thus produced rises to the surface 
in large bubbles when the mud is stirred up with a stick. This 
gas is, of course, not pure methane, but contains, besides, some 
carbonic acid and nitrogen-the former being produced by decom¬ 
position, the latter from the air dissolved in the water. 

Methane is set free in immense quantities in some coal-mines, 
owing to partial decomposition of the coal. A mixture of this 
methane with atmospheric air, which is often produced in such 
mines, explodes when ignited, even more violently than a mixture 
of hydrogen and air. Two volumes of hydrogen, requiring for its 
combustion one volume of oxygen, produce two volumes of water- 
gas ; but two volumes of methane, requiring four volumes of 
oxygen, produce, on ignition, two volumes of carbonic acid, and - 
four volumes of water-gas : — 

CH 4 + 20 2 = co 2 + 2H 2 0. 

2 vols. 4 vols. 2 vols. 4 vols. 

With a falling barometer—*>. with reduced air pressure, large 
quantities of this methane pass from cavities in the coal into the 
workings of the mines, and, mixing with the air, produce an ex¬ 
plosive mixture (fire-damp), which is readily ignited by the lamp 
of the miner, and then produces those fearful explosions by which 
so many lives are annually lost. 

The safety lamp constructed by Davy to avoid these explosions 
consists of a small oil-lamp completely surrounded by wire gauze, 
and its action depends upon the fact that an ignited and explosive 
caseous mixture when passing through wire gauze has its tempera- 

o 


Methane, or Marsh Gas. 305 

ture lowered below its point of ignition by the conductivity of the 
metal for heat, and, therefore, does not burn on the other side of 
the gauze. Various unforeseen difficulties have, however, prevented 
its general application, and it is often rendered useless by the 
miner opening the gauze casing for purposes of his own. 

Methane is produced, not only by slow putrefaction, but also 
by the destructive distillation of coal, and is therefoie one of the 
chief constituents of coal-gas. The well-known lightness of coal- 
gas and its use for filling balloons are due to this methane, and to 
the hydrogen which it contains. Other carbon compounds also 
yield methane when strongly heated. Thus, the vapour of ordinary 
alcohol: C 2 H fi O, when driven through a red-hot iron tube, yields 
considerable quantities of methane, mixed however with carbonic 
acid :— 

2 C 3 H 6 0 = 3 ch 4 + C 0 2 . 

Methane may be obtained fairly pure by strongly heating an 
intimate mixture of dry sodium acetate and caustic soda, or better 
soda-lime 1 in a tube of hard glass. The products are then 
methane arid sodium carbonate :— 

|CH 3 

(CO-ONa + NaOH = CH 4 + CO(ONa) 3 . 

Sodium acetate 

Chemically pure methane can only be obtained from an organic 
compound called zinc methyl : Zn(CH 3 ) 2 , and which is distin¬ 
guished by its great tendency to unite with oxygen. This body is a 
colourless, volatile liquid, which catches fire at once when exposed 
to the air and decomposes water with almost explosive violence. 
If zinc methyl is boiled in a glass flask, out of contact with air, and 
its vapour then passed into water, it is at once decomposed into 
methane and zinc hydrate, the latter separating out as a white, 
gelatinous solid, or remaining dissolved as zinc chloride, if hydro¬ 
chloric acid has been previously added to the water :— 

Zn(CH 3 ) 2 + 2H 2 0 = 2CH 4 + Zn(OH) 2 . 

Zinc methyl 

Methane is a perfectly indifferent compound, and unites with 
no other substance without decomposition. Sulphuric, nitric, or 
phosphoric acid are without action on it, and even the powerful 
oxidizing mixture of potassium dichromate and sulphuric acid, as 

1 Quick-lime which has been slaked with caustic soda solution instead of 
with water.—E d. 


X 



3o 6 Text-Book of Inorganic Chemistry. 

well as caustic potash, leave it entirely unchanged. Chlorine is 
without action upon it in the dark, but if a mixture of the two 
•gases is led into vessels exposed to direct sunlight, one or more 
atoms of hydrogen, according to the volume of the chlorine, are 
abstracted from the methane to form hydrochloric acid, while the 
place of the hydrogen atoms in the methane molecule is taken by 
an equal number of atoms of chlorine. The end product of the 
action of chlorine on methane contains no hydrogen, but as many 
atoms of chlorine as were originally present of hydrogen. 

The composition of these different chlorine substitution-products 


of methane is as follows :— 

Methane (Methyl hydride) . . . . CH 4 

Monochlormethane (Methyl chloride) . . CH 3 C 1 

Dichlormethane (Methylene chloride) . . CH 2 CI 2 

Trichlormethane (Chloroform) . . . . CHC 1 3 

Tetrachlormethane (Carbon tetrachloride) . CC 1 4 


These compounds belong to organic chemistry. 

Ethylene, also known as Olefiant gas, because it forms an oily 
liquid with chlorine, contains in its molecule the same quantity of 
hydrogen as methane but twice as much carbon, and has, there¬ 
fore, the composition : C 2 H 4 . Ethylene is a colourless gas, of 
peculiar odour, slightly soluble in water, and condensed to a liquid 
under considerable pressure and at a low temperature. Its specific 
gravity is 0-968, and its molecular weight 0*968 x 28*88 = 28*00, 
corresponding, therefore, to the formula : C 2 H 4 . It is easily in¬ 
flammable and burns with a bright luminous flame. 

Ethylene is easily obtained from common alcohol. This com¬ 
pound, which, at a low red heat, breaks up into carbonic acid and 
methane, is converted by the abstraction of the elements of water 
into ethylene. Concentrated sulphuric acid is best employed for 
this purpose. Four parts by weight of. the acid are gradually 
added to one part of strong alcohol in a large flask provided with 
a delivery tube. The mixture is then gently heated, and as soon 
as its temperature has risen to about 150°, a copious evolution of 
ethylene begins. To free the gas from alcohol carried over with 
it, it is passed through a wash-bottle containing water, and to 
remove the sulphurous anhydride, produced towards the end of the 
operation, a second wash-bottle containing caustic potash is used. 


Ethylene or Olefiant Gas. 307 

The process is represented by the following equation, the sulphuric 
acid remaining unchanged 

c 2 h 6 o = C 2 H 4 + H 2 0 . 

H ethylene and chlorine are mixed in a large vessel and ex- 
p osed to diffuse daylight (in direct sunlight the mixture would 
explode), the two gases unite to produce a colourless, heavy oil, of 
pleasant ethereal odour, and having the composition : C„HXL— 
ethylene dichlonde. Ethylene also unites directly with bromine 
producing ethylene dibromide : C 2 H 4 Br 2 , a liquid with similar 
properties to the chloride. The corresponding iodide -CHI is 
a crystalline solid, but less stable. 2 4 25 

Ethylene behaves, therefore, like a dyad element, and unites 
without change with two atoms of chlorine, &c., just like copper 
zinc, and other metals. ’ 

If a glass cylinder is half filled with ethylene over water, then 
an equal volume of chlorine added and a burning taper quickly 
brought near the open end of the jar, the mixture catches fire, 
and burns with an exceedingly smoky flame. The products are 
hydrochloric acid and carbon. Even when excess of chlorine is 
present no chloride of carbon is formed, a proof that the attraction 
of carbon for chlorine is not sufficient to cause the elements to 
unite directly with one another. 

Since the molecule of ethylene contains two atoms of carbon, 
instead of one, like methane, it requires a larger quantity of oxygen 
to burn it than this gas. Two volumes of ethylene require & six 
volumes of oxygen for complete combustion, and produce four 
volumes of carbonic acid, and four of water-gas :_ 

+ 30 2 = 2 C 0 2 + 2 H 2 0 . 

6 vols. 4 vols. 4 vols. 

Ethylene and oxygen mixed in this proportion— i.e. 1:3, explode 
violently when ignited, and easily burst a thin glass flask which has 
been carefully enveloped in towels. 


c 2 h 4 

2 vols. 


At first sight it seems strange that ethylene should burn with 
a flame so much more luminous than methane, although the latter 
gas contains half as much carbon and the same quantity of hydro¬ 
gen ; and connected with this point is the question, Upon what does 
the luminosity of a flame depend ? 

Experiment teaches us that under normal conditions only those 
flames are luminous which contain a glowing solid body. The 


308 Text-Book of Inorganic Chemistry. 

scarcely luminous hydrogen and oxy-hydrogen flames give out 
an intense light if a solid substance—^, platinum, lime, &c.—is 
strongly heated in them. 

In the same way the flame of burning ethylene contains a solid 
substance— i.e. carbon. And the presence of carbon in the flame 
may be readily proved by holding a cold, white porcelain dish in 
it. Those parts of the dish which come into contact with the 
inner parts of the flame receive a black deposit of soot, just as 
burning arsine deposits black arsenic on a piece of cold porcelain 
(p. 232). 

The constitution of a luminous flame is not, however, so simple 
as might be thought from the above statements. Three parts may 
be distinguished in every luminous flame : an innermost 
and cold portion, an intermediate and luminous, and a 
less luminous and hot external mantle, as is illustrated 
in fig. 57. The gas which streams out of the burner at 
A does not at once enter into combustion, because no 
oxygen is present with which it can combine. The inner 
cone, a'a, consists therefore of unburnt gas, and is cold. 
This may be readily proved by stretching a fine platinum 
wire across the flame, which will become red-hot at the 
two sides of the flame but remain dark, because cold, in 
the centre. 

The oxygen which rapidly diffuses from the exterior 
to the interior portions of the flame not only causes the 
gases to burn in the exterior mantle, but, by the high 
temperature produced during this combustion, decom¬ 
poses the unburnt gases in the interior of the flame. The decom¬ 
position which is thus brought about is the same as that which 
ethylene suffers when passed through a red-hot tube out of con¬ 
tact with the air. It is then changed into methane and carbon : 
C H 4 = CH 4 + C, and it is this carbon which is so liberated, and 
which is heated in the interior mantle of the flam ef e, g (fig. 57) 
by the external burning gases, that gives the flame its luminosity. 
Ethylene is not, therefore, directly burnt in the flame, but rather the 
products into which it has been changed—viz. methane and carbon. 

The same process goes on when a candle burns. A burning 
body brought near to the wick produces the same changes in the 
fat or wax as destructive distillation, and forms combustible gases 
which catch fire. The heat of the flame then continually melts 



Fig- 57 . 




Carbon Tetrachloride. 


309 


the solid fat at its base, which is sucked up by the porous wick to 
the flame to be decomposed and burnt. It is not, however, the fat 
itself which burns, but the gaseous products into which it has been 
decomposed by the heat of the flame. The further course of the 
combustion then resembles that of ethylene. 

A luminous flame may easily be made non-luminous if the 
combustible gases are mixed with enough oxygen to completely 
oxidize the whole of the carbon. This can either be done by 
blowing air or oxygen into the luminous flame (as in the blowpipe), 
or by previously mixing the gas with air. In the Bunsen burner a 
mixture of gas and air is obtained in the following manner. The 
gas is allowed to issue from a small opening over which is screwed 
a brass tube, with two openings at the base, to admit the air. The 
gas and air then mix so thoroughly in ascending the tube, that the 
mixture burns at the top of the tube without smoke or luminosity. 

The temperature of a flame is sometimes judged from the in¬ 
tensity of the light which it emits, and although this may to a 
certain extent be true for solid bodies (iron is hotter when white- 
hot than when red-hot, and when brightly glowing than when feebly 
glowing), in the case of burning gases luminosity and temperature 
are often inversely proportional to one another. 

The non-luminous flame of coal gas burning in a Bunsen burner 
is much hotter than the luminous flame of the same gas burnt m 
the usual manner. In the latter case the interior of the flame 
contains unburnt carbon, but in the former all the constituents of 
the gas are at once burnt. For the same reason the external non- 
luminous mantle d, b, c, of the gas flame (fig. 57 ) is much hotter 
than the interior luminous portion. 


CARBON TETRACHLORIDE. 

Of the numerous compounds of carbon and chlorine which are 
more closely allied to organic than inorganic compounds, we shall 
here only briefly allude to the most important and at the same time 

the simplest. 

Composition : CC 1 4 . 

Carbon tetrachloride is a colourless liquid sinking in water but 
not mixing with it. Its specific gravity is i'6, and its boiling 



310 Text-Book of Inorganic Chemistry. 

point 77° The liquid possesses a pleasant, ethereal odour, and 
when breathed acts as an anaesthetic like chloroform. 

Carbon tetrachloride cannot be obtained by the direct union of 
carbon and chlorine, but is readily produced from carbon-disulphide 
when this liquid is mixed with an excess of dry chlorine and 
exposed to the light. A quicker method of preparing it consists in 
passing chlorine which has bubbled through warm carbon di¬ 
sulphide through a red-hot porcelain tube, or by treating this last- 
named liquid with antimony pentachloride in the presence of 
chlorine. In all cases the carbon disulphide is converted into 
carbon tetrachloride and disulphur dichloride. These mixed 
liquids are then digested with caustic soda, which decomposes the 
latter, and the carbon tetrachloride then distils over with water 
vapour on heating. It has already been mentioned (p. 306) that 
carbon tetrachloride can also be obtained from methane. 

Carbon tetrachloride is a very stable compound ; it is distin¬ 
guished from most other inorganic chlorides by the fact that when 
boiled with aqueous potash it remains unchanged. Its alcoholic 
solution is not decomposed when a similar solution of silver nitrate 
is added to it —silver chloride is not precipitated. 

If gaseous carbon tetrachloride is led through a red-hot tube it 
is decomposed into chlorine, and a liquid chloride of carbon con¬ 
taining only one half as much chlorine. The molecule of this com¬ 
pound does not, however, consist of one atom of carbon and two 
of chlorine, but contains two atoms of carbon united with four atoms 
of chlorine, and has, therefore, a similar composition to the tetra¬ 
chloride as ethylene has to methane 

Methane . . . CH 4 Tetrachlormethane . » CC 1 4 

Ethylene . . . C 2 H 4 Tetrachlorethylene . . C 2 C 1 4 

This tetrachlorethylene, like ethylene, unites with two atoms 
of chlorine, and yields a compound corresponding to ethylene 
dichloride :—- 

Ethylene dichloride : C 2 H 4 C 1 2 . 

Tetrachlorethylene dichloride (carbon hexachloride) : C 2 C 1 6 . 

Carbonic hexachloride is a volatile crystalline solid with an 
odour resembling that of camphor, and which, like camphor, easily 
sublimes on the walls of the vessels in which it is kept. 

Cai'bon Tetrabromide : CBr 4 , closely resembles the tetrachloride 
in its preparation and properties. 


Cyanogen. 


3“ 


CYANOGEN. 

Cyanogen is the name given to a compound consisting of one 
atom of carbon and one atom of nitrogen, which is so similar in its 
properties to an element, and especially to the halogens, that it 
would certainly have been considered as such if it were not pos¬ 
sible to produce it from its constituents, and to again decompose it 
into them. 

Cyanogen was discovered by Gay-Lussac in 1815 ; its name is. 
derived from icvaveos , blue, because it is a constituent of Prussian 
blue, which can be prepared from it. 

Composition : CN. 

The molecule of cyanogen consists of two atoms, 1 just as the 
molecule of hydrogen also contains two atoms, and the composition., 
of its molecule is therefore expressed by the formula : CN - CN, or 
(CN) 2 . Cyanogen may also be still more briefly expressed by the 
symbol : Cy. 

Cyanogen is a colourless gas of a peculiar odour, resembling 
hydrocyanic acid, and with a specific gravity of i*8. This gives a 
molecular weight of r8x 28*88 = 52, corresponding to the formula 
(CN) 2 . At a low temperature, and under a strong pressure, it may 
be condensed to a colourless liquid, boiling at — 20°, or even to a 
crystalline solid. Water dissolves about 4^ times its volume ; 
alcohol considerably more. It is easily inflammable, and burns 
with a beautiful flame of a peach-blossom colour, which, as neither 
glowing carbon nor nitrogen possess this colour, must be peculiar 
to glowing cyanogen. 

Cyanogen is prepared by heating mercuric cyanide. This white 
crystalline compound, which is easily soluble in water, is obtained 
by dissolving mercuric oxide in excess of aqueous hydrocyanic 
acid, and then crystallizes out on evaporating the clear solution. 

1 Strictly, it is, of course, incorrect to speak of an atom of cyanogen, since it 
is not, like the atoms of the e’ements, chemically indivisible. It is, however, 
convenient to extend the word atom to the compound radicals of organic 
and inorganic chemistry, which play the part of elementary substances 
instead of coining a new word such as semi-molecule . When we speak of an 
atom of cyanogen, ammonium, amidogen, or ethyl, we do not mean an indivi¬ 
sible quantity of the substance, but such a quantity which exactly corresponds 
to the elementary atoms and can displace them in chemical compounds. 


312 Text-Book of Inorganic Chemistry . 

After being freed from traces of water which it contains, it breaks 
up when heated to low redness into mercury and cyanogen, just 
as mercuric oxide decomposes under similar circumstances into 
mercury and oxygen. The gas can be collected over mercury. A 
portion of the cyanogen remains behind in the tube as an amor¬ 
phous brown substance called paracyanogen , which has the same 
composition as cyanogen but a higher molecular weight, and of 
which it is therefore a polymeric compound. 

No compound of cyanogen with oxygen or sulphur is known in 
the free state, and it can only be brought into combination with these 
substances indirectly. With the halogens it combines at once as 
soon as they liberate it from its compounds with hydrogen (hydro¬ 
cyanic acid) or with the metals (potassium or mercuric cyanide.) 

Potassium and sodium burn in cyanogen like they burn in 
chlorine, and unite directly with the gas, producing potassium or 
sodium cyanide. These compounds are distinguished by their 
stability; they can be heated to redness, out of contact with the 
air, without decomposition. 

Although carbon and nitrogen never unite directly with one 
another when heated together alone, they may be made to combine 
in the presence of a third substance which readily unites with 
cyanogen— eg. potassium. Potassium carbonate and carbon, when 
heated to bright redness, form carbonic oxide and potassium, and 
if nitrogen is led over the glowing mixture, a compound of potas¬ 
sium and cyanogen is produced, called potassium cyanide : KCy, 
a white solid, soluble in water like potassium chloride, and, like 
this salt, yielding, when decomposed bp sulphuric acid, an acid— 
hydrocyanic acid—corresponding to hydrochloric acid. 

The same conditions necessary for the production of potassium 
cyanide are also present when nitrogenous organic substances, 
such as blood, hair, horn, are mixed with potassium carbonate and 
heated. Under these circumstances, the nitrogen and carbon of 
the organic substances unite with the potassium to form potassium 
cyanide. This process is that by means of which cyanogen com¬ 
pounds are manufactured on a large scale. The first compound 
always prepared is the so-called yellow prussiate , or potassium 
ferrocyanide, which may be considered as a double cyanide of 
potassium and iron : 4KCy,Fe ,/ Cy 2 , and which will be more fully 
discussed under the compounds of iron. It is also produced when 
an aqueous solution of potassium cyanide acts upon powdered 
ferrous sulphide ; a portion of the potassium cyanide is then decom- 


Hych'ocyanic Acid. 313 

posed, and forms potassium sulphide and ferrous cyanide. From 
its solution in water it crystallizes in large yellow tablets. 

Potassium ferrocyanide forms the starting-point for the prepara¬ 
tion of all the other cyanogen compounds— e.g. hydrocyanic acid, 
potassium cyanide, potassium cyanate, &c. 


HYDROCYANIC ACID. (Prussic Acid.) 

Composition : HCN or HCy. 

This compound, well known on account of its highly poisonous 
character, is, when completely anhydrous, a colourless, mobile, 
volatile liquid, miscible with water in all proportions. It possesses 
a strong stupefying odour, and has a specific gravity of o'j. Its 
boiling-point is 27 0 , and at - 15 0 it solidifies to a crystalline solid. 
It is easily combustible. 

Anhydrous hydrocyanic acid may be prepared in the same way 
as hydrochloric acid by pouring concentrated sulphuric acid on 
to potassium cyanide, but the reaction is much too violent, and, on 
account of the poisonous nature of the gas, far too dangerous to be 
employed. It is much better to first prepare the aqueous acid and 
then to abstract water from it. For the same reason it is better 
to employ for the preparation of the dilute acid not potassium 
cyanide, but its compound with ferrous cyanide— i.e. potassium 
ferro-cyanide. 

To obtain the dilute aqueous acid, 10 parts of powdered potas¬ 
sium ferrocyanide are covered with a mixture of 6 parts of sulphuric 
acid and 30 parts of water in a tubulated retort with its neck in¬ 
clined upwards. The end of the neck of the retort is connected 
by a cork and a bent tube with a condenser, the other end of which 
dips in a receiver carefully cooled in ice, and in which it is best to 
place a little water at the commencement. The mixture in the 
retort is then gradually brought to boiling, when a portion of the 
water-vapour condenses in the neck of the retort, while the volatile 
hydrocyanic acid passes over to the well-cooled receiver. The 
whole apparatus should be placed in a good upward draught, and 
every precaution taken, not only in the preparation of this highly 
poisonous volatile compound, but also in all experiments made with 
it afterwards. 


3 



314 Text-Book of Inorganic Chemistry. 

In order to prepare the anhydrous compound from its aqueous 
solution, a flask which is about one-third filled with the latter is 
placed in ice and small pieces of fused calcium chloride gradually 
dropped in, which unites with the water and partially liquefies. 
On no account must the aqueous acid be poured on to the calcium 
chloride, as the heat produced by the union of the latter with the 
water would drive off large quantities of gaseous hydrocyanic acid. 
After a few hours, the flask is attached to a long condenser and a 
receiver immersed in a freezing mixture, and then gently heated 
on a water bath. Anhydrous hydrocyanic acid then distils over as 
a clear, transparent liquid. 

Aqueous, and still more anhydrous hydrocyanic acid, undergoes 
a change when kept for some time in closed vessels, with the 
separation of a brown amorphous substance, known as parahydro- 
cyanic acid , and probably a polymeric compound of hydrocyanic 
acid. The decomposition is usually accompanied with an evolution 
of gas, which sometimes exerts sufficient pressure to destroy the 
vessels in which it is preserved. It has been found that the addi¬ 
tion of a few-drops of hydrochloric acid materially diminishes this 
spontaneous decomposition. 

Hydrocyanic acid is one of the most powerful poisons known. 
A few drops of the strong aqueous acid produce almost instanta¬ 
neous death. Its vapour is also highly poisonous when respired ; 
a pigeon dies in about half a minute when breathing air mixed with 
the vapour from anhydrous hydrocyanic acid, even when the tempe¬ 
rature of the liquid is as low as o°. 

Very largely diluted with water, hydrocyanic acid is a valuable 
medicine. The strength prescribed by the ‘ British Pharmacopoeia ’ 
is a 2 per cent, solution ; a 3 per cent, solution is also sometimes 
used under the name Scheele’s prussic acid. Bitter almonds, and 
the kernels of cherries, plums, &c., as well as the leaves of the 
common or cherry laurel, contain a crystalline organic compound 
of complex composition, called amygdaline. When these sub¬ 
stances are crushed and moistened with water, the amygdaline 
comes into contact with a ferment (emulsine) which they also con¬ 
tain, and is then decomposed as sugar is decomposed by yeast. 
One of the products of this decomposition is hydrocyanic acid, 
which, when the water is afterwards distilled, passes over with the 
water-vapour into the receiver. The preparations known as bitter- 
almond-water and laurel-water are obtained in this way. 

Hydrocyanic acid is so weak an acid that it does not even red- 


Hydrocyanic Acid\ 315 

den blue litmus paper, and is expelled by carbonic acid from its 
compounds with the alkalies. For this reason, potassium cyanide 
when exposed to the air smells strongly of hydrocyanic acid. 

In many of its reactions, it shows a similarity with the hydrogen 
compounds of the halogens—for example, when mixed with a solution 
of silver nitrate it gives a white, curdy precipitate of silver cyanide, 
which, like silver chloride, is insoluble in nitric acid but dissolves 
in ammonia. Silver chloride may, however, be at once distin¬ 
guished from silver cyanide if heated; the former compound 
simply melts, but the latter is decomposed into silver and cyanogen 
gas. As potassium chloride unites with platinic chloride to form 
a stable crystalline double compound, so also potassium cyanide 
unites with nearly all the metallic cyanides, and produces from 
these insoluble substances soluble crystalline double cyanides, of 
which the stable potassium ferrocyanide may be taken as an 
example. 

Chlorine easily decomposes hydrocyanic acid, uniting with its 
two constituents to form cyanogen chloride and hydrochloric 
acid. 

In order to detect hydrocyanic acid in a liquid, advantage is 
taken of the readiness with which Prussian blue may be formed 
from it. Prussian blue may be considered as a double compound 
of ferric and ferrous cyanides, Fe /// Cy 3 and Fe^Cy^ and although 
insoluble in water and dilute acids, is not produced when a mixture 
of a ferrous and ferric salt is added to the aqueous acid. If, how¬ 
ever, this mixture is made alkaline with caustic potash, potassium 
ferrocyanide, and a mixture of ferrous and ferric hydrates are pro¬ 
duced, and this, when acidulated with hydrochloric acid, forms 
Prussian blue due to the union of the potassium ferrocyanide with 
the ferric chloride. 

Cyanogen is a monad radical, and is to be considered as a com¬ 
pound of triad nitrogen with dyad carbon. According to this, hydro¬ 
cyanic acid would be a compound resembling ammonia in which 
two atoms of hydrogen are displaced by one atom of dyad carbon, 

(C" 

and would then possess the formula : N j H And the production 

of silver cyanide from hydrocyanic acid would be represented by 
the equation :— 

n{£ + N 0 2 • OAg = N j^ g + NCVOH. 


31 6 Text-Book of Inorganic Chemistry . 

At present we only know one kind of hydrocyanic acid, but it is 
possible, and indeed probable, that a second isomeric compound 
exists, containing tetrad carbon as the grouping element. In this 
compound, three of the four bonds of the carbon atom would be 
satisfied with the atom of nitrogen, and one with the atom of 

hydrogen, thus : C j ^ This hydrocyanic acid would probably 

be only slightly, if at all, poisonous. Cyanides are known which 
from their peculiar chemical behaviour appear to contain two 
modifications of cyanogen, from which the existence of a second 
modification of hydrocyanic acid may be deduced. 

Possible, but not probable, is the existence of a third isomeric 
hydrocyanic acid containing pentad nitrogen as grouping element 
saturated with one atom of tetrad carbon and one of hydrogen, 
and of which the composition would be expressed by the formula : 


CYANIC ACID. 

Composition : CONH. 

Although cyanogen cannot unite directly with oxygen as it does 
with chlorine, potassium cyanide can, under favourable circum¬ 
stances, easily take up an atom of oxygen and become oxidized to 
potassium cyanate. But the compound called cyanic acid does not 
possess a similar composition to chloric acid ; such a compound is, 
in fact, as much unknown as compounds of cyanogen corresponding 
to perchloric or chlorous acid, or any of the oxides of chlorine. 
Cyanogen appears in fact to unite with oxygen and hydrogen, or a 
metal in only one proportion—viz. as cyanic acid : CONH. 

Potassium cyanate, although similar in compositio 7 i to potassium 
hypochlorite, ClOK, has by no means the same constitution. In 
the hypochlorite the atom of oxygen unites together the chlorine 
and potassium, but in the cyanate the oxygen atom is not united 
to the cyanogen as a whole, but to the atom of dyad carbon, form¬ 
ing carbonic oxide, which then plays the part of the carbon in 
potassium cyanide. The relation of potassium cyanide to potas¬ 
sium cyanate is best expressed by the following rational for¬ 
mulae :— 



Cyanic Acid. 


317 


Potassium cyanate 

whence it follows that the constitution of 


Potassium cyanide 



~C\” 

cyanic acid is N 



The change of potassium cyanide into the cyanate, which takes 
place so readily, consists therefore of the oxidation of the atom of 
carbon and in its change from the dyad to the tetrad state. 

Potassium cyanate, a crystalline salt, soluble in alcohol and 
water, is readily obtained by dropping lead oxide into fused potas¬ 
sium cyanide when the oxide is reduced to metallic lead. The 
salt will be more fully described under the potassium compounds. 

Cyanic acid cannot, however, be set free from this salt by a 
stronger acid, because whenever the acid comes into contact with 
water it is at once decomposed into carbonic acid and ammonia :— 



C 0 2 + NH 3 . 


Cyanic acid may be obtained by heating a polymeric compound 
—cyanuric acid, C 3 N 3 0 3 H 3 —which is a white solid, obtained from 
urea. Cyanuric acid is not itself volatile, but when heated breaks 
up into cyanic acid, which distils over as a colourless liquid, with 
an odour resembling acetic acid. If kept below o° it may be pre¬ 
served for some time unchanged, but at a few degrees above this 
temperature the liquid becomes heated, and changes into a white 
solid, resembling porcelain. This compound— cyamelide —is iso¬ 
meric with cyanic acid, and is again converted into it on heating. 

The compounds of cyanogen, especially those belonging to 
organic chemistry, are remarkable for the readiness with which 
they change into isomeric and polymeric compounds. 


SULPHOCYANIC ACID. 
Composition : CSNH. 


Just as fused potassium cyanide unites readily with oxygen, 
even that of the air, so also it possesses a strong attraction for sul¬ 
phur. Sulphur dropped into fused potassium cyanide readily com¬ 
bines with it, producing potassium sulphocyanate : CSNK, which 



318 Text-Book of Inorganic Chemistry . 

from its analogy with potassium cyanate may be considered to 
possess the same constitution : N The corresponding hydro¬ 

gen compound, sulphocyanic acid : CSNH, is more stable than 
cyanic acid ; an aqueous solution of it may be prepared by dis¬ 
tilling potassium sulphocyanate with dilute sulphuric acid. The 
anhydrous acid is best obtained by passing dry sulphuretted hydro¬ 
gen'over powdered mercuric sulphocyanate. It then condenses in 
the receiver as a colourless oily liquid, with a piercing odour, also 
resembling acetic acid, but when kept for a short time soon de¬ 
composes into hydrocyanic acid and a yellow crystalline solid 
nearly insoluble in water —persuiphocyamc acid : C 2 H 2 N 2 S 3 . 

Sulphocyanic acid or any soluble sulphocyanate can be at once 
recognized by the bright red colour which it gives with ferric 
chloride. This reaction is quite as delicate and distinct as the 
blue precipitate produced from potassium ferrocyanide and ferric 
chloride. 


Chlorides of Cyanogen. 

We are acquainted with two polymeric compounds of cyanogen 
and chlorine, one a volatile liquid and the other a crystalline solid. 
These compounds appear to stand in the same relation to one 
another as cyanic acid to cyanuric acid. 

Liquid Cyanogen Chloride : CNC 1 , is a colourless liquid, 
boiling at 12°, solidifying at — 7 0 , and with an exceedingly sharp, 
tear-producing odour. It is very poisonous. This compound is 
obtained by the action of chlorine on dilute hydrocyanic acid or 
on metallic cyanides in the presence of water. The simplest 
method of preparation is to lead a slow stream of chlorine into a 
solution of potassium cyanide with twice its weight of water kept 
cooled in ice. The solution is placed in a retort, which it only fills 
to about one-fourth, and the neck of the retort is connected with a 
tube filled partly with bright copper turnings (to absorb the excess 
of chlorine), and partly with calcium chloride (to retain moisture). 
This tube communicates with a U-tube placed in a freezing mixture 
in which the cyanogen chloride condenses. A portion of the volatile 
compound distils over as soon as it is formed ; the residue is after¬ 
wards obtained by heating the retort to 50°. 

Pure liquid cyanogen chloride may be preserved unchanged in 
sealed tubes, but when impure it soon changes into a solid poly- 


Coal-Gas. 


319 


meric modification. Water and alcohol dissolve considerable 
quantities without decomposition, and it can be again expelled 
from these solutions on heating. Alkalies decompose it. Liquid 
cyanogen chloride unites chemically with the chlorides of other 
elements— e.g. of antimony, boron, titanium. 

Solid Cyanogen Chloride , probably C 3 N 3 C 1 3 , is produced by the 
action of chlorine on anhydrous hydrocyanic acid in bright sunlight, 
when it separates out in colourless lustrous crystals. Like the 
liquid chloride, it has a piercing odour. It melts at 145 0 , boils at 
190°, and is decomposed on boiling with alkalies, or even water, 
into cyanuric and hydrochloric acids. 

Cyanogen Bromide'. CNBr, and Cyanogen Iodide'. CNI, are 
volatile crystalline solids, and are produced by treating mercuric 
cyanide or potassium cyanide with bromine or iodine respectively. 
Both possess a similar penetrating odour to that of the two 
chlorides. Cyanogen iodide is often contained in commercial 
iodine. 


COAL-GAS. 1 

Coal-gas is a complex mixture of various gases produced by 
the destructive or dry distillation of coal. The coal when distilled 
also yields a number of liquid and solid products, which are con¬ 
tained in coal-tar —a thick, oily, black liquid ; and in the ammo- 
7iiacal-liquor —an aqueous liquid, smelling strongly of ammonia and 
ammonium sulphide. The apparatus used in the manufacture of 
coal-gas consists of the retorts in which the coal is heated, the 
condensers for removing the coal-tar and ammoniacal-liquor, the 
purifiers for separating gaseous impurities, and gasometers for 
storing the gas. 

The retorts are made of fire-clay, and so that their cross section 
has a D-shape; they are closed at one end and are arranged in 
the furnace in a horizontal position, with their open ends projecting 
slightly outwards. After the charge of coal has been introduced, 
the open ends are closed by an iron plate, which can be bolted on, 
and from which a vertical iron tube passes upwards. This iron 

1 By the Editor. 



320 


Text-Book of Inorganic Chemistry. 


tube, which conveys away all the products of distillation from the 
retorts, dips under the surface of water contained in a large iron 
tube running in a horizontal position above the retorts and called 
the hydraulic main . Most of the tar and much of the ammonia is 
here condensed, and when the liquid reaches a certain height 
it runs off through an opening into the tar-well —a large reservoir 
receiving all the liquid products of distillation. The gas next 
passes through a series of upright iron tubes, called atmospheric 
condensers, which are freely exposed to the air, and so cool the gas 
and condense a further quantity of liquid products. But in order 
to completely remove these substances, as well as some of the 
gaseous impurities, the gas is passed through the scrubbers , or tall 
vessels filled with fragments of coke, and kept continually moist 
by a descending spray of cold water. The gas enters at the base 
of the scrubber and escapes near the top. It now only remains to 
remove traces of gaseous impurities, especially sulphuretted 
hydrogen, carbon disulphide, and other sulphur compounds, 
which impart an unpleasant odour to the gas, and which, when 
burnt, produce sulphurous and sulphuric acids. This final 
purification is effected in the purifiers , which consist of large 
chambers containing thin layers of dry slaked-lime or of ferric 
hydrate on shelves. The lime then absorbs the sulphuretted 
hydrogen and other sulphur compounds, together with any car¬ 
bonic acid which may be present, and is converted into spent- 
lime —a dirty-green amorphous mass of various compounds of 
calcium, with a very offensive odour. If ferric hydrate is employed 
it is converted by sulphuretted hydrogen into ferrous sulphide and 
free sulphur, and the former compound is partially oxidized on 
afterwards exposing the mass to the air. But when the quantity 
of free sulphur becomes large, the ferric hydrate can no longer 
effect the purification, and is then sent to the sulphuric acid works 
to burn out the sulphur and again obtain pure ferric oxide. Ferric 
hydrate does not remove carbonic acid, and it is, therefore, better to 
mix it with a certain proportion of slaked-lime in order to get rid 
of this impurity. The presence of sulphuretted hydrogen in coal- 
gas may be detected by allowing a jet of the gas to play on a piece 
of paper impregnated with a solution of lead acetate. If the paper 
does not become blackened after a few seconds, the gas may be 
said to be free from this impurity. This test does not, however, 
detect the presence of carbon disulphide or other sulphur com¬ 
pounds. 


Coal-Gas. 


321 

The gases contained in coal-gas may be divided into three 
classes :— 

(i.) Diluents , which have practically no illuminating power, but 
which are of special value when the gas is used for heating 
purposes. These consist of hydrogen , marsh gas , and car¬ 
bonic oxide, and constitute from 85 to 90 percent, by volume 
of the entire gas. 

(ii.) Illummants , to which the illuminating power of the gas is 
due, and which consist of hydrocarbons containing more 
carbon in their molecule than marsh gas. They are some¬ 
times called heavy hydrocarbons. These illuminants are 
principally ethylene (C 2 H 4 ), propylene (C 5 H 0 ), butylene (C 4 H 9 ), 
_ acetylene (C 2 H 2 ), and benzene (C e H„). 

(iii.) Impurities , which consist of sulphuretted hydrogen , carbon 
disulphide and other sulphur compounds, together with traces 
of carbonic acid , oxygen , and nitrogen , the two latter being 
derived from the air. 

The composition of coal-gas depends upon the temperature at 
which the coal is distilled and upon the kind of coal used. The 
gas also varies in composition at different stages of the distillation. 
If the temperature is high, the heavy hydrocarbons are decomposed 
and a large volume of gas of low illuminating power is obtained. 
And if cannel coal is employed the gas contains a much larger 
proportion of these heavy hydrocarbons than that obtained from 
ordinary bituminous coal. The following table gives the compo¬ 
sition of two samples of coal-gas. No. 1 is ordinary London gas, 
made from a mixture of bituminous coal and a small quantity of 
cannel ; No. 2 is gas made entirely from cannel coal and supplied 
to the Houses of Parliament :— 


Marsh Gas 

Hydrogen 

Carbonic Oxide 

Heavy Hydrocarbons 

Nitrogen 

Oxygen 

Carbonic Acid 


No. 


35'89 
50-68 
3'98 
4-08 

4’93 
0-26 
traces 
99-82 


No. 2 


41-88 

4172 

478 

8-72 
271 
none 
none 
100-oi 


The greater part of the nitrogen which the coal contains is 
converted into ammonia during the distillation. An aqueous 

Y 

















322 


Text-Book of Inorganic Chemistry. 

solution of this ammonia constitutes the ammoniacal-liquor which 
collects in the tar-well. The ammonia is contained in this liquid 
partly in the free state, and partly in combination as ammonium 
carbonate, sulphide, cyanide, &c. This ammoniacal-liquor is now 
the chief source of ammonia and its compounds. 

The coal-tar which is also found in the tar-well is a highly 
complex mixture of a large number of organic compounds. Among 
these may be mentioned the liquids-benzene and its homologues, 
aniline and phenol (carbolic acid); and among the solids—paraffin, 
napthalene, and anthracene. When the tar is distilled the more 
volatile liquids pass over and produce a combustible mixture of 
various compounds, called coal-tar naphtha. Coal-tar is used for 
a variety of purposes, some of which depend upon the antiseptic 
properties of the phenol which it contains, but by far the largest 
quantities are consumed in the manufacture of the beautiful coal- 
tar or aniline colours. 

After the distillation is finished a considerable portion of the 
carbon present in the coal remains behind in the retorts as a grey 
porous mass, containing also the whole of the ash of the coal. This 
constitutes coke. At the same time, owing to the decomposition of 
some of the heavy hydrocarbons during distillation, a layer of ex¬ 
tremely hard and compact carbon collects on the upper surface of 
the retorts. This form of carbon is called gas-carbon ; it is used 
in the construction of some galvanic batteries and for highly refrac¬ 
tory crucibles. 


TITANIUM. 

Chemical Symbol'. Ti.— Atomic Weight'. 48. 

This element, which in many respects is closely allied to silicon, 
is also always found in nature, combined with oxygen. 

Titanium dioxide or titanic anhydride : Ti 0 2 , either free or 
combined with bases, is the form in which the element always 
occurs. Free titanic anhydride is found in three varieties : rutile , 
anatase , and brookite, which, although chemically the same, have 
different crystalline and other physical characteristics. Rutile 
crystallizes according to the tetragonal system, anatase and brookite 
according to the rhombic, and titanic anhydride is therefore trimor- 
phous. Of the other naturally occurring compounds of titanium, 



Titanium. 


323 

titanic i?'on, probably ferrous titanateand titanite or sphene , calcium 
titanate and silicate, are the most important. 

Free titanium is obtained in the following manner. Two iron 
boats, one containing dry potassium fluotitanate : K 2 TiF 6 , and the 
other metallic sodium, are brought into a tube of hard glass, which 
is then filled with pure hydrogen. The boat containing the fluo¬ 
titanate is first heated to redness, then that containing the sodium, 
so that the vapour of the metal is carried by the hydrogen over 
the hot titanium salt. By this process sodium and potassium 
fluorides are produced and free titanium. If the two former com¬ 
pounds are afterwards extracted by hot water the titanium remains 
behind as a dark-grey amorphous powder, resembling iron which 
has been reduced in hydrogen. The element, in this pulverulent 
form, burns brilliantly when heated in oxygen or the air, forming 
titanic anhydride : Ti 0 2 . It dissolves easily in hydrochloric acid, 
with evolution of hydrogen, probably producing a solution of 
titanium dichloride : TiCl 2 . It burns when heated in chlorine, 
giving titanium tetrachloride : TiCl 4 . 

The material used for the preparation of titanium compounds 
is nearly always rutile. This substance, like quartz, is insoluble 
in hydrochloric acid, but may be brought into solution either by 
fusing with caustic potash or potassium carbonate, or with acid 
potassium sulphate, or, finally, by heating with concentrated sul¬ 
phuric acid until the greater part of the liquid has been volatilized. 

Titanic Anhydride : Ti 0 2 .—The titanium sulphate produced 
by either of the last two methods dissolves in cold water, and from 
this solution titanic acid is precipitated in white flocculent masses 
on the addition of alkalies or alkaline carbonates, and is not re¬ 
dissolved by an excess of these substances. The dilute aqueous 
solution of titanium sulphate is decomposed on boiling, and deposits 
titanic acid as a white powder, which is only slightly soluble in 
concentrated sulphuric acid. 

Potassium titanate, obtained by fusing powdered rutile with 
three times its weight of potassium carbonate, does not dissolve 
when the fused mass is afterwards treated with water, but forms 
acid potassium titanate, which is insoluble in water, and which can 
therefore be purified by washing. This acid salt dissolves, how¬ 
ever, in cold hydrochloric acid, and when the solution is diluted 
with water and heated to boiling, titanic acid is precipitated as a 
white powder. 


y 2 


324 Text-Book of Inorganic Chemistry . 

This form of titanic acid is scarcely soluble in water containing 
hydrochloric acid or ammonium chloride, but when washed with 
pure water passes in a milky state through the filter. 

The titanic acid precipitated by boiling aqueous solutions con¬ 
taining hydrochloric or sulphuric acid always contains traces of 
the acid mixed with it, from which, however, it may be completely 
freed by digesting with ammonia, which has no action on the titanic 

The white precipitate produced by ammonia in a cold hydro¬ 
chloric acid solution of titanic acid is said to have the composition : 
Ti(OH) 4 . This compound is easily soluble in dilute acids, but is 
converted on warming into an acid containing less water—meta- 
titanic acid, of the probable composition : TiO(OH) 2 , which is on y 
slightly soluble in acids. The titanic acid precipitated by boiling 
a sulphuric acid solution is insoluble in acids except concentiated 
sulphuric acid. 

If metallic zinc is added to a hydrochloric acid solution of 
titanic acid, the nascent hydrogen which is produced reduces the 
titanic acid to the hydrated sesquioxide. In consequence of this 
reduction the liquid becomes of a violet-blue colour, and after a 
time, if the solution was not too dilute, a dark blue or violet blue 
precipitate separates out, which again becomes white when exposed 
to the air from absorption of oxygen. 

Titanium Sesquioxide : Ti 2 0 3 , is obtained in the anhydrous state 
by glowing titanic anhydride in a stream of hydrogen. It is a black 
powder insoluble in nitric or hydrochloric acid, but dissolving in 
sulphuric acid to form a violet solution. When strongly glowed in 
the air it becomes white, being reconverted into titanic anhydride. 

Titanium Tetrachloride: TiCl 4 .—This compound, which 
closely resembles silicon tetrachloride, except that it is less volatile, 
is obtained, like silicon chloride, by heating an intimate mixture 
of rutile and charcoal in chlorine. It then distils over into a cooled 
receiver as a heavy liquid, usually with a yellowish tinge due to 
traces of ferric chloride. When purified by redistillation its 
specific gravity is 176, and boiling point 136°. It fumes strongly 
in moist air and has a vapour density of 6*8, corresponding to the 
molecular formula : TiCl 4 . Water at once decomposes it, with 
considerable evolution of heat, into hydrochloric and titanic acids. 
Titanium chloride absorbs considerable quantities of ammonia 


Molybdenn m, 325 

yielding a reddish-brown powder of the composition : TiCl 4 ,4NH 3 , 
which is easily decomposed by moist air. 

Titanium Tetrafluoride : TiF 4 , is a colourless, fuming liquid, 
obtained by heating a mixture of titanic anhydride and fluor-spar 
with concentrated sulphuric acid. 

Fluotitanic Acid: H 2 TiF fi , which is analogous in composition 
to fluosilicic acid, is obtained by dissolving titanic acid in hydro¬ 
fluoric acid. If the aqueous solution is neutralized with potassium 
carbonate, or if potassium titanate is dissolved in hydrofluoric acid, 
or, finally, if rutile is fused with acid potassium fluoride, potassium 
fluotitanate : K 2 TiF 6 , is easily obtained in the crystalline form. 

Titanium belongs to those elements which unite with nitrogen 
to form compounds stable at a red heat. If titanium chloride is 
saturated with dry ammonia, and the resulting solid : TiCl 4 , 4NH3 
heated to redness in gaseous ammonia, titanium nitride : Ti 3 N 4 , 
remains behind as a copper-coloured mass, with metallic lustre, 
and which evolves ammonia when fused with solid caustic potash. 

In the slag from blast furnaces, where iron ores containing 
titanium have been used, bright-red, lustrous crystals in cubes are 
often found, and sometimes in considerable quantities. They consist 
of a chemical compound of titanium nitride and cyanide. 


MOLYBDENUM. 

Chemical Symbol'. Ho.—Atomic Weight'. 96. 

Molybdenum occurs chiefly in nature as molybdenite , the di¬ 
sulphide : MoS„ and more rarely as wulfenite , lead molybdate: 
Mo 0 2 - 0 2 Pb. Molybdenite, which closely resembles graphite in 
its appearance, is a soft mineral, feeling greasy to the touch, and 
easily marking paper. It is, however, heavier than graphite, and 
is further distinguished by giving an olive-green streak on porcelain, 
instead of a grey one like graphite. 

Molybdenum is readily obtained by heating one of its chlorides 
or oxides in hydrogen, when the walls of the glass tube used for 
the reduction often become covered with a brilliant, metallic-like 
mirror of the free substance. In its physical properties, molyb¬ 
denum far more closely resembles the metals than the non-metals. 



326 Text-Book of Inorganic Chemistry. 

It is of a silver-white colour, with a strong metallic lustre, very 
hard, melts with extreme difficulty, and is possibly infusible when 
quite pure. Its specific gravity is 8‘6. In the air it remains un¬ 
changed, and only combines with oxygen when heated to low red¬ 
ness. Hydrochloric acid or dilute sulphuric acid has no action on 
it. Concentrated sulphuric acid converts it into a brown mass. 
Aqua regia easily dissolves it, and nitric acid converts it into one 
of the oxides. 

For the preparation of molybdenum compounds from molyb¬ 
denite, the mineral is powdered, and placed in an open, inclined 
crucible, which is then surrounded with pieces of glowing char¬ 
coal and heated to low redness, with frequent stirring, as long as 
sulphurous anhydride is evolved. Crude and impure molybdic 
anhydride then remains behind as a yellow powder. This com¬ 
pound is insoluble in water, but dissolves in aqueous ammonia, 
leaving the impurities and the unoxidized molybdenite behind. 
A further quantity of molybdenum may generally be obtained from 
the residue by roasting again as before. On evaporating down 
the ammoniacal solution ammonium molybdate crystallizes out, and 
from a solution of this salt, cautious addition of dilute hydrochloric 
or nitric acid precipitates molybdic acid as a white powder, scarcely 
soluble in water, but easily soluble in an excess of acid. 

Molybdenum forms several compounds with oxygen. We dis¬ 
tinguish, for example, molybdenum sesquioxide : Mo 2 0 3 , obtained 
by reduction from molybdic anhydride as a black powder insoluble 
in acids, also molybdenum dioxide : Mo0 2 , which remains behind as 
a red crystalline mass when ammonium molybdate is heated in a 
closed crucible, and 

iviolyhdic Anhydride : Mo0 3 , the best known oxide. Molyb¬ 
dic anhydride or molybdenum trioxide remains as delicate white 
crystals, yellowish when hot, on heating ammonium molybdate in 
an open crucible. It melts at a red heat, and sublimes in small, 
thin, glistening scales. Acids do not dissolve it, but with alkalies 
it forms soluble crystalline compounds—the molybdates. 

iviolyhdic Acid : Mo0 2 (OH) 2 , is a dibasic acid, and separates 
from a solution of any of its salts on careful addition of dilute 
hydrochloric acid as a white crystalline powder, easily soluble in 
excess of acid. If zinc is added to a hydrochloric acid solution of 
molybdic acid, a blue colour is first produced which soon becomes 


Tungsten. 3 2 7 

1 green, and finally dark-brown. These changes in colour are due 
to the reduction of the molybdic acid by the nascent hydrogen. 

Among the salts of molybdic acid that of ammonium is of 
considerable importance in analytical chemistry ; it is used for 
the detection and estimation of phosphoric acid in acid liquids 
containing other compounds. If such a liquid containing phos¬ 
phoric acid—for example, the aqueous or acid extract of a soil or 
manure—is mixed with an excess of a solution of ammonium 
molybdate strongly acidified with nitric acid and then warmed, 
a yellow crystalline precipitate of the composition : iiMo0 3 , 
P0(0NH 4 ) 3 ,6H 2 0, and containing the whole of the phosphoric 
acid, is thrown down. This precipitate is insoluble in water or 
dilute acids, but easily soluble in ammonia. If this yellow compound 
is boiled with aqua regia, the ammonia is destroyed, and on allow¬ 
ing the solution to evaporate at the ordinary temperature, phospho- 
molybdic acid : i iMo0 3 ,PO(OH) 3 ,I2H 2 0, crystallizes out in yellow 
prisms. The acid is easily soluble in water, and produces, with 
acid solutions of ammonium salts or nitrogenous organic bases, 
yellow compounds which are insoluble in dilute acids. 

Ammonium molybdate, when mixed with an acid solution of 
arsenic acid or an arsenate, produces a similar yellow insoluble salt. 

Several chlorides of molybdenum are known as : MoCb, 
MoC1 3 , MoC1 4 , MoC 1 5 , and MoC1 6 ; besides these two oxychlorides 
of hexad molybdenum also exist, viz. : MoOC1 4 and Mo0 2 C1 2 . 


TUNGSTEN. 

Chemical Symbol \ W .—Atomic Weight'. i 84 * 

This element resembles molybdenum in many respects, but is 
distinguished from it by the high specific gravity of the free element 
and of its compounds. 

Tungsten never occurs free in nature, and its compounds aie 
only found in small quantities. The best known of the tungsten 
minerals are wolfram (ferrous and manganous tungstate) and schee - 
lite (calcium tungstate), the former of a dark grey colour with a 
specific gravity of 7 - 5 , the latter consisting of white crystals also 
remarkable for its high specific gravity. 



328 Text-Book of Inorganic Chemistry . 

Tungsten, prepared by the reduction of tungstic anhydride in a 
stream of hydrogen at a bright red heat, forms small grey crystal¬ 
line particles, which under the burnisher become lustrous like 
iron. Its specific gravity is about 19 , and thus nearly approaches 
that of gold. It is only difficultly fusible, especially in large 
quantities, is brittle, and so hard that it scratches glass. At the 
ordinary temperature it remains unchanged in the air ; but when 
it is heated in the pulverulent form it burns easily, and pro¬ 
duces tungstic anhydride. Hot nitric acid oxidizes it to the same 
product. 

Of the oxides of tungsten, the dioxide, W0 2 , is known as a 
brown powder insoluble in hydrochloric or nitric acid. More im¬ 
portant than this is the trioxide— 

Tungstic Anhydride : WO s , which remains behind as a 
yellow powder, insoluble in water and acids, when powdered 
wolfram is heated with nitric acid, while the iron and manganese go 
into solution. It may also be obtained by acting on scheelite with 
nitric or hydrochloric acid. This yellow tungstic anhydride dis¬ 
solves readily in alkalies (including ammonia), forming soluble 
crystalline salts—the tungstates. From a cold solution of these 
salts dilute acids precipitate white tungstic acid of the composition : 
WO(OH) 4 , which, when dried over sulphuric acid, passes into the 
dibasic acid : W0 2 (0H) 2 , and when heated to the boiling point of 
water is converted into yellow tungstic anhydride. 

If tungstic acid is precipitated from a solution of a tungstate by 
hydrochloric acid, and then a piece of zinc added to the acid liquid, 
a beautiful dark blue colour is produced, due to a soluble oxide of 
uncertain composition, and which, when exposed to the air, again 
oxidizes to tungstic acid. 

The normal tungstates are salts of the dibasic acid : W0 2 (0H) . 
Others, however, possess a more complex composition ; for exam¬ 
ple, the metatungstates of which the sodium salt has the formula : 
Na 2 W 4 0 13 + ioH 2 0. The common sodium tungstate of commerce 
is still more complex, and has the composition : Na 10 W 12 O 41 . 

Corresponding to phosphomolybdic acid is a phosphotungstic 
acid of similar composition. And a remarkable double compound 
of tungstic and silicic acids is known —silicotungstic acid , the 
soluble and crystalline compounds of which are obtained by boil¬ 
ing acid alkaline tungstates with gelatinous silicic acid. 

Like molybdenum, tungsten unites in several proportions with 


Vanadntm . 


329 


chlorine. The hexachloride: \VC1 0 , corresponding to tungstic 
anhydride, forms a dark violet mass, melting at 275 ° to a red 
liquid, and boiling at 346 °. It is decomposed by water into 
tungstic and hydrochloric acids. This compound may be obtained 
by moderately heating free tungsten in a stream of pure chlorine. 
Chlorides of composition : WC1 5 , WC1 4 , WC1 2 , are also known. 
Corresponding to tungsten hexachloride are two oxychlorides : 
WOCl 4 and W0 2 C1 2 , both solid crystalline substances. 

Tungsten, when alloyed, in small quantities, with iron, increases 
its hardness; steel containing about 5 per cent, of tungsten- 
known as tungsten-steel —is distinguished by its great hardness. 


VANADIUM. 

Chemical Symbol \ V.— Atomic Weight: 51 '3- 

Vanadium belongs to the rarer elements, although it is widely 
distributed in minute quantities ; it occurs especially in certain 
ores of iron. Vanadic acid is usually prepared from these iron 
ores, the mineral vanadinite— lead vanadate—being too rare to 
serve for the preparation of vanadium compounds. Considerable 
quantities of vanadium compounds are now obtained in England 
from a mineral called mottramite , a double vanadate of lead and 
copper, occurring in Lancashire. 

Vanadium, obtained by glowingthedichloride in pure hydrogen, 
is a grey powder of specific gravity 5 - 5 , which remains unchanged 
in the air, and can be raised to redness out of contact with air with¬ 
out melting, but when heated in the air burns to form vanadic 
anhydride. It also unites readily with nitrogen at a red heat. It 
is insoluble in hydrochloric acid, but dissolves easily in nitric acid 
to a blue solution. Fused with caustic soda, hydrogen is evolved, 
and sodium vanadate produced. 

Of the oxides of vanadium the best known is :— 

Vanadic Anhydride : V 2 0 5 , analogous in composition to phos¬ 
phoric anhydride, and, like this, forming an acid, but little known 
in the free state, which is tribasic. Vanadic anhydride, obtained 
by gently roasting ammonium vanadate, is a reddish powder, which 
melts when heated to redness, and crystallizes on cooling. It is 
slightly soluble in water, giving a solution of a yellow colour, but 



330 Text-Book of Inorganic Chemistry . 

dissolves easily in acids or alkalies, forming with the latter soluble 
crystalline vanadates. 

Metallic zinc brought into contact with a hydrochloric acid 
solution of vanadic acid produces first a blue colour, which after¬ 
wards changes to green. 

The extraction of vanadic anhydride from iron ores containing 
vanadium may be carried out in the following manner. The ores 
are finely powdered, mixed with nitre and heated. The mass, 
which contains potassium vanadate, is then extracted with water, 
nearly neutralized with nitric acid and then precipitated with 
barium chloride. Insoluble barium vanadate is thus produced, 
which is washed with water, and then decomposed by boiling with 
sulphuric acid. The acid liquid is then separated from the barium 
sulphate, neutralized with ammonia, concentrated, and solid pieces 
of ammonia chloride added, when ammonium vanadate, which is 
soluble in water, but insoluble in ammonium chloride, gradually 
separates out. This salt, when washed with ammonium chloride 
solution, dried and gently heated in an open crucible, leaves vanadic 
anhydride behind. 


NIOBIUM AND TANTALUM. 

These rare elements are but little known either in the free state 
or in combination with other substances. Their compounds are 
most nearly related to those of vanadium. They both form anhy¬ 
drides, corresponding to vanadic anhydride, and having therefore 
the composition : Nb 2 0 5 niobic anhydride and Ta 2 0 5 tantalic an¬ 
hydride, in which the elements are pentads. 

Niobium and tantalum nearly always occur associated with one 
another in nature ; they are found in the minerals colunibite and 
tantalite , which are essentially ferrous niobate and tantalate, in 
samarskite , a niobate and tantalate of the metals of the yttrium and 
cerium groups, and in some other still rarer minerals, from which 
the oxides are best extracted by fusion with acid potassium sulphate. 

Niobic and tantalic anhydrides are white powders, yellow when 
hot, and insoluble in water, acids, or alkalies. They are principally 
distinguished from one another by their very different specific 
gravity. 



33i 


THE METALS. 

The classification of the elements into the two great divisions, 
metals and non-metals, dates from a time when a number of 
physical properties were ascribed to the metals, which were 
thought to be peculiar to them and which were considered suit¬ 
able for their distinction from the non-metals. Such characteristic 
and peculiar properties of the metals were, for example, their 
lustre (whence the term ‘metallic lustre’), their conductivity for. 
heat and electricity, a high specific gravity (exceeding 6), opacity to 
light, and others. 

We now know that not one of these physical characteristics 
belongs exclusively to the metals. Iodine, tellurium, and graphite 
possess metallic lustre. Tellurium conducts heat, and graphite, as 
well as selenium (in a less degree) conduct electricity. On the 
other hand, metals are known— e.g. potassium and sodium—of 
which the specific gravity is not only less than 6, but which are 
even lighter than water. 

The exact meaning, therefore, of the words metal and non-metal 
cannot be sharply defined—it is impossible to say what properties 
are exclusively peculiar to the former and what to the latter, just as 
it is impossible to exactly classify any series of natural substances. 

This, however, by no means prevents us from grouping to¬ 
gether, according to artificial rules, certain classes of natural 
phenomena and substances which possess some similarity with 
one another; and by so doing we not only obtain a clearer con¬ 
ception of the entire subject, but also make it easier for the 
student. 

It cannot be denied that a more comprehensive view of the 
more than sixty elements known to modern chemistry is obtained 
if we classify those resembling one another into certain groups. 
This has already been done with the elements of the sulphur and 
nitrogen groups, and with the halogens. For the same reason, it is 
therefore advisable to retain the division of the whole of the ele¬ 
ments into non-metals and metals; and since the application of 


332 


Text-Book of Inorganic Chemistry. 

certain physical properties for a sharp classification has proved 
useless, we must seek for other peculiarities which shall serve to 
separate oxygen, hydrogen, carbon, the halogens, the members of 
the sulphur and nitrogen groups, &c., from potassium, aluminium, 
zinc, lead, silver, &c., when writing or speaking of chemical facts. 

An artificial classification of the elements cannot be obtained 
from their physical properties, and we must therefore base our 
divisions on their chemical behaviour. For this purpose we can 
use the chemical character of their compounds with oxygen. 

The larger number of the elements when united with oxygen 
produce oxides of a basic character , the other elements produce 
almost exclusively oxides with acid properties. To the latter class 
belong those elements which we have as yet discussed , and which 
are called non-metals. Those elements of which the oxides usually 
possess basic prope?'ties are the metals. 

That this division of the elements into non-metals and metals 
is an artificial and not a natural classification, and that it cannot 
therefore be consistently carried out, is proved by the following 
considerations:—(i.) Oxygen belongs to neither class ; but is 
reckoned with the non-metals, (ii.) Hydrogen often plays the part 
of a metal, and its oxide, water, behaves sometimes as a base and 
sometimes as an acid, (iii.) Antimony, although included under 
the non-metals, forms a basic oxide with oxygen—antimonous 
oxide, (iv.) The metals manganese and chromium form oxides 
of acid properties—manganic and chromic anhydrides—as well as 
others of basic properties, (v.) The basic oxides of many metals— 
e.g. lead oxide, zinc oxide—play the part of acids in some com¬ 
pounds. Lead oxide, which combines with nitric acid to form 
lead nitrate, also unites with potash and forms a salt—potassium 
plumbite. In the former compound it plays the part of a base, in 
the latter that of an acid. 

It is clear, therefore, that the division of the elements accord¬ 
ingly as they produce acid or basic oxides cannot, like every other 
artificial classification, be carried out in every particular ; it is 
only suited to separate certain elements which differ from one 
another in certain important particulars. 

The Chemical Constitution of Salts. 

With the definition of a metal is connected that of a salt, since 
all salts of inorganic chemistry (if we. consider ammonium as a 


Constitution of Salts. 333 

metal) are metallic compounds ; and a short, accurate definition 
of a salt is scarcely easier than that of a metal. 

On p. 78 it has been stated that salts are produced by the 
union of acids with bases, which tells us how they are produced, 
but not of what they actually consist. 

Salts are binary compounds consisting of an electro-positive 
constituent— e.g. a metal, and an electro-negative 1 constituent— 
e.g. a halogen or acid radical. They are usually decomposed by 
electrolysis into these two constituents—the former separating at 
the negative, the latter at the positive pole. Common salt (sodium 
chloride) when fused or dissolved in water is decomposed by an 
electric current into chlorine and sodium, of which the former 
goes to the positive pole and the latter to the negative. In the 
presence of water the sodium is decomposed, so that instead of 
sodium we obtain hydrogen and sodium hydrate at the negative 
pole. 

Salts of this kind contain a metal in direct union with a nega¬ 
tive element, the halogen. They are called haloid-salts because 
only the halogens and the compound substances resembling them, 
cyanogen, sulphocyanogen, &c., but not the other non-metals, can 
thus unite directly with the metals to form salts. 

Another class of salts, called amphid-salts , consist also of an 
electro-positive and electro-negative constituent, which, however, 
are not directly united with one another, but are combined 
through the intervention of one or more atoms of oxygen or sul¬ 
phur.^ If the uniting element is oxygen, the compound is called an 
oxysalt, if sulphur a sulphosalt. 

A haloid-salt may be converted into an oxysalt by the intro¬ 
duction of an atom of oxygen—for example, in the conversion of 
sodium chloride into sodium hypochlorite : NaCl + O = NaOCl, in 
which the sodium and chlorine are not now directly united with 
one another, but through the atom of oxygen, common to both. 

The negative portion, and sometimes the positive portion also, 
of the oxysalts usually consists of a compound radical. Thus, 
potassium nitrate : N0 2 ‘0K, contains the monad radical nitric 

peroxide (nitryl): N0 2 , sodium sulphate : S0 2 1 Q]sj a contains the 
dyad radical sulphurous anhydride (sulphuryl) : S0 2 , and sodium 

1 Using the obsolete but convenient nomenclature of the electro-chemical 
theory.—E d. 


334 


Text-Book of Inorganic Chemistry. 


rONa 

phosphate : PO- ONa the triad radical : PO, which has not yet 
l ONa 


been isolated. And these three acid radicals are all united to their 
respective metals—the positive portion of the salt, by one, two, or 
three atoms of oxygen respectively . 1 Two metals may also be 
united with one another by oxygen, as in potassium plumbite : 

Pk jOK * n l ea -d is the negative element, and the potas¬ 

sium the positive. 

As an example of the sulphosalts may be mentioned the com¬ 
pound produced by the combination of carbon disulphide with 

I SK 

potassium sulphide, potassium sulphocarbonate : CS j corre- 

f OK 

sponding to ordinary potassium carbonate: CO|q^ in which 

compounds the respective oxygen and sulphur atoms may displace 
one another. For although a compound of the composition: 
(OK 

CO | gj^'is at present unknown, the previously described (p. 168 ) 

f ONa. 

sodium thiosulphate: S 0 2 |g^ a is an example of these little 

known inorganic compounds. In organic chemistry they are much 
more numerously represented. 

Besides these two elements, oxygen and sulphur, it also 
appears that in certain cases the element fluorine can assume a 
dyad character, and so serve to unite the negative and positive 
portions of a salt (p. 136 ). 


We also distinguish salts called normal acid , and basic salts. 
Monobasic acids— i.e. those containing only one atom of hydrogen 
displaceable by one atom of a monad metal, eg. nitric acid : 
N0 2 -0H—can only form normal salts with monacid bases— i.e. 
those containing only one atom of hydrogen displaceable by an 
acid radical, eg. caustic potash : KOH. 

Acid salts can only be produced from polybasic acids— i.e. those 
which contain more than one atom of displaceable hydrogen, as 
sulphuric or phosphoric acid, when a portion of their hydrogen is 
displaced by a metal (comp. p. 56 ). In the same manner, basic 


salts are formed from polyacid bases— eg. lead hydrate : Pb 


JOH 

10H 


1 The electrolysis of such an oxysalt has already been explained (p. 79 ).—Ed. 


Constitution of Salts. 


335 


OH 


and ferric hydrate : Fe , OH when a portion only of their hydro- 
l OH 

gen is displaced by an acid radical or by a haloid. The following 
are some examples of basic salts : — 

0-N0 o 
OH 

Cl 
OH 

ono 2 

OH 
OH 

o 2 so 2 

(OH 
(Cl 2 
lOH 


Basic lead nitrate 

. Pbj 

Basic lead chloride 

. . Pbj 

Basic ferric nitrate 

' - Fe { 

Basic ferric sulphate . 

. . Fe( 

Basic ferric chloride . 

• Fe i 


Many compounds which are called basic compounds, but which 
do not contain the acid radical or haloid and the metal in equiva¬ 
lent proportions, are best considered as mixtures of a normal salt 
and a metallic hydrate. Thus the basic antimonous chloride, 
referred to on p. 257 , and sometimes represented by the formula : 


Sb q contains varying quantities of oxygen and chlorine united 

with the same quantity of antimony, and is, therefore, better con¬ 
sidered as a mixture of antimonous chloride : SbCl 3 , and anti¬ 
monous oxide: Sb 2 O s . 

(OH 

A triacid base— e.g. aluminium hydrate : AV OH requires three 

OH 


molecules of monobasic acid in order to produce a normal salt, 
while at the same time three molecules of water are also produced : 
A 1 ( 0 H ) 3 + 3 N 0 2 0 H=A 10 3 (N 0 2 ) 3 + 3H 2 0; and in the same way 
a molecule of the same base will require only one molecule of a 
tribasic acid: Al(OH ) 3 + PO(OH ) 3 = A10 3 -P0 + 3H 2 0. If, how¬ 
ever, we have a dibasic acid united to the same triacid base, we 
shall require three molecules of the former to two of the latter. 
Thus, normal aluminium sulphate is produced in this way :— 

2 Al(OH ) 3 + 3S0 2 (OH ) 2 = A1 2 0 6 (S0 2 ) 3 + 6H 2 0. 


Normal salts may also be produced by the union of two mole¬ 
cules of a dibasic acid with one of a triacid, and one of a monacid 



336 • Text-Book of Inorganic Chemistry. 

base. Common alum: k 0 3 ) 2SO *’ for exam P le > is a normal 

sulphate of the two metals aluminium and potassium, and is thus 
formed :— 

Al(OH), + KOH 4- 2S0 2 (OH) 2 = ^jaSO, + 4H 2 0. 

From these examples it will be seen that a salt may be defined 
either as an acid in which the hydrogen is displaced by a metal, or 
as a base in which its hydrogen is displaced by an acid radical, or 
its hydroxyl by a haloid element or some similar group of elements. 

The normal salts often have no action on vegetable colours — 
i.e. are said to be neutral, and hence this class of salts is sometimes 
called neutral salts. But some normal salts have a strong alkaline 
or acid reaction, and even some acid salts react occasionally alka¬ 
line, while, on the other hand, some acid salts possess a neutral 
reaction. Red litmus paper is strongly blued by normal potassium 

(OK 

carbonate: CO ■ OK an d by normal potassium phosphate : PO-lOK 
[OK (OK 

(OK 

and even faintly by monacid potassium phosphate : PO] OK On 

(OH 

(ONa 

the other hand, acid sodium carbonate : CO 1 qjj reacts neutral. 

The cause of this is that the strongly basic properties of potassium 
hydrate are not neutralized by the weak carbonic acid, and that 
phosphoric acid, although one of the strongest acids, does not 
completely neutralize the potash in the monacid phosphate. 

Double Salts. —Compounds in which the chemical affinities of 
the combining atoms are completely satisfied are called saturated 
compounds. Such, for example, are potassium chloride : KC1, and 
platinum tetrachloride : PtCl 4 . And it might be thought that these 
two compounds could not be capable of any further combination. 
But in spite of this they both unite to form a double, and very stable 
salt of the composition : 2 KCl,PtCl 4 . And a large number of other 
similar double salts are also known. 

In most of these cases where perfectly saturated compounds 
can unite with one another chemically and form new compounds, 
we are unacquainted with the forces which cause this union. Such 
compounds have been called molecular compounds — i.e. compounds 


Constitution of Salts. 337 

produced by the chemical union of molecules. This expression, 
of course, explains nothing, but simply states that the molecules of 
saturated compounds are capable of uniting with one another to 
produce new compounds. 

The constitution of the haloid double salts may be expressed in 
the following manner. From a consideration of the composition 
of numerous double fluorides, it appears that two atoms of monad 
fluorine can coalesce to form what plays the part of a dyad atom. 
Thus in the compound KF, HF (comp. p. 136 ), we can consider 
that the two monad atoms of fluorine unite to form one dyad atom 
of double the atomic weight, which maybe indicated by the symbol 
¥. This double compound then becomes KFH, analogous to 
KOH. Just in the same manner the double chloride of potassium 
and platinum, which may be written : 2 KCl,(PtCl 2 )Cl 2 , may be 
considered as two atoms of potassium united with the dyad radical 
PtCl 2 by the other four atoms of monad chlorine, which have 
coalesced to form two dyad atoms, thus :_ 

k 2 ci 2> (Ptcyci, = £§}ptcv 

In the same manner the double iodide of potassium and mercury: 
2 KI, Hgl 2 may be written: ^|| Hg, and the double cyanide of po¬ 
tassium and silver, KCyAg. The constitution of the potassium 
ferrocyanides, as well as other similar compounds, may be inter¬ 
preted in the same way. 

The oxysalts also unite together to form double salts. Zinc 
sulphate, for example, unites with potassium sulphate to form the 
compound : S0 2 0 2 Zn + S0 2 (OK ) 2 + 6H 2 0, which is produced from 
zinc sulphate by the substitution of a molecule of water of consti¬ 
tution, by potassium sulphate. 


SPECTRUM ANALYSIS. 

It has been previously mentioned (p. 308 ), that a non-luminous 
flame becomes luminous when a solid body is raised to incande¬ 
scence in it, and then emits white light. But if metallic compounds 
are introduced into such a flame and there volatilize, for which the 
metallic chlorides are best adapted, the heated vapour of the 
metal imparts a colour to the flame, peculiar to each. Thus 

z 




338 Text-Book of Inorganic Chemistry . 

compounds of sodium colour the flame an intense yellow, those of 
potassium violet, and those of lithium crimson. But owing to the 
similarity of many of these colours when viewed by the naked eye, 
they are not always suitable for directly determining which metal 
may or may not be present. It is often difficult to say by mere 
observation whether a given flame is coloured red by strontium or 
calcium, or whether a flame coloured yellow by sodium, also con¬ 
tains potassium, since the bright yellow sodium flame can com¬ 
pletely mask the violet colour due to potassium compounds, even 
when the latter substance is present in considerable quantities. 

But something quite different is observed if, instead of simply 
looking at a flame with the naked eye, we allow the light to fall 
through a narrow slit on to a triangular prism of glass and then 
magnify the bent rays of light by a telescope. Light of different 
colours in passing through the prism is differently bent, and 
ordinary white light gives a band of colours extending from red to 
violet—a so-called continuous spectrum. The light from a glow¬ 
ing gas is broken up into its constituent parts in the same way . 
the spectrum in this case will not, however, be a continuous one 
but will consist of bright lines separated by more or less dark 
spaces. And these bright lines always appear in those parts of 
the spectrum to which they belong; thus the red lines will always 
be in the red, the blue in the blue portion, and so on—each gaseous 
element having its own peculiar spectrum. 

The coloured table of spectra forming the frontispiece shows 
firstly the continuous spectrum of sunlight and then the spectra of 
ten metals, which differ much from one another, both with regard 
to the number and to the position of their lines. Sodium, lithium, 
thallium, and indium possess at the most two lines each, while 
Ccesium, rubidium, strontium, calcium, and barium have much 
more complex spectra. Caesium, strontium, and indium have all 
blue lines, but their position, when compared with the same scale, 
is different. As is shown in the table, the blue strontium line is 
most to the left, then follow the two blue lines of caesium more to 
the right, and still further to the right is the blue indium line. 

A substance which is supposed to contain thallium, and which 
like thallium gives a green line but in a different position—either 
more to the right or to the left—certainly does not contain this 
metal. Such an observation would rather point to the existence of 
a new (hitherto unknown) element; for, as previously remarked, 
the position of the bright lines in the spectrum of any one element 


Spectrum Analysis. 339 

is always the same under all conditions. In this way the new metals, 
rubidium, caesium, thallium, indium, and gallium, have been dis¬ 
covered by the aid of spectrum analysis, and it is extremely probable 
that others will be added to this list. And although, as previously 
intimated, potassium compounds, even when present in consider¬ 
able quantity, cannot be recognized by the naked eye in a flame in 
which sodium is volatilized at the same time, if we allow the light 
to fall through a slit and decompose it by the prism, the lines of 
both metals are easily seen in the spectrum. 

The apparatus for spectrum analysis—called a spectroscope— 
as devised by Bunsen and Kirchhoff, to whom this delicate method 


Fig. 58. 

of chemical analysis is due, is shown in one of its simplest forms 
in fig. 58 . On an iron stand are fastened three brass tubes, and 
between these is placed the glass prism. At the outside end of 
the tube A is a narrow slit, which can be widened or contracted 
at will, and through which the light from the heated substance 
passes on to the glass prism ; the spectrum so produced is then 
observed through the telescope B. The third tube, C, contains a 
small reduced photographic scale, illuminated by a small gas flame 









340 Text-Book of Inorganic Chemistry. 

placed in front of it, and an image of which is thus reflected from 
one face of the prism into the telescope b. This serves to fix the 
position of any line in a given spectrum. All extraneous light must, 
of course, be excluded from the prism by covering it with a black¬ 
ened box or cloth. 

A little practice with the spectroscope soon enables one to 
determine the presence or absence of any metal in a given sub¬ 
stance. . 

Those metals which are only volatilized with difficulty [eg. 
iron), and which therefore do not become gaseous in the gas flame, 
are investigated by allowing powerful electric sparks to pass between 
poles of the metal, or between small carbon poles saturated with 
a solution of its chloride. Gases yield spectra in the same way. 
Hydrogen, when heated by the passage of electric sparks, appears 
bright red, and its spectrum consists essentially of a bright red, a 
green, and a blue line. 

The results of spectrum analysis have also made it possible to 
determine the composition of the sun and other fixed stars. And 
in no heavenly body has an element been discovered which is un¬ 
known on the earth. 

The great delicacy of spectrum analysis makes it especially 
valuable for chemical purposes. Most minute quantities of various 
metals, which it was previously quite impossible to detect, can by 
recognized with the greatest certainty by spectrum analysis. It 
has been calculated from experiments that one hundred-thousandth 
part of a milligramme of lithium chloride, and even one three- 
millionth part of a milligramme of sodium chloride may be easily 
detected by their spectra. 


CLASSIFICATION OF THE METALS. 

It has been repeatedly attempted to discover some physical or 
chemical properties of the metals which might serve to classify 
them into a series of divisions and sub-divisions. But, however 
desirable such a classification may be in simplifying the study of 
this large group of elements, all such attempts have only strength¬ 
ened the conviction that the metals cannot be exactly divided 
upon any natural system into a number of distinct classes, and 
that, if such a classification is made, it must be according to some 



34 i 


Classification of the Metals. 

artificial rules. The properties of the oxides and sulphides of the 
metals may serve as the basis of such an artificial classification. 

One class of the metals forms oxides, or rather hydrates, which 
are easily soluble in water, the solutions of which have a powerful 
alkaline reaction; these hydrates are called the alkalies, and the 
corresponding metals the metals of the alkalies. The following 
five metals belong to this class—viz. potassium , sodium , lithium , 
rubidium , and ccesium , and the hypothetical compound metal 
ammonium. 

A second class also yields soluble oxides, though far less 
soluble than the alkalies. These oxides also react faintly alkaline, 
but are at the same time allied to the oxides of the next—the so- 
called earths —and their corresponding metals are therefore called 
the metals of the alkaline earths. In this class the metals calcium , 
strontium , barium , and magnesium are included. The last-named 
metal, of which the oxide is only very slightly soluble in water, forms 
a connecting link with the next group. 

The third group, of which the oxides are completely insoluble 
in water, and are called earths, are the metals of the earths. The 
chief representative of this group is aluminium , with which are 
associated the rarer metals, beryllium , indium , gallium , yttrium , 
erbium , cerium , lanthanum , didymium , and zirconium , as well as 
some still rarer elements. 

The sulphides of this group (except indium) are not produced on 
the addition of sulphuretted hydrogen or an alkaline sulphide to a 
solution of one of their salts. The precipitate produced on adding 
sodium sulphide or ammonium sulphide to their solutions usually 
consists of the insoluble hydrate, while sulphuretted hydrogen is 
given off:— 

A1 2 C1 6 + 3 Na 2 S + 6H 2 0 = 2A1(0H) 3 + 6NaCl + 3 H 2 S. 

A fourth class includes those metals of which the sulphides are 
insoluble in water, but mostly soluble in dilute acids. These com¬ 
pounds are precipitated on adding an alkaline sulphide, such as 
ammonium sulphide, to a solution of one of their salts. They are 
not, however, produced when sulphuretted hydrogen is led through 
an acid solution of their salts. The metals of this group are : iron ., 
manganese , chromium , uranium , cobalt , nickel, and zinc. 

The fifth group includes those metals of which the sulphides are 
insoluble both in water and in dilute acids, and are therefore 
precipitated when sulphuretted hydrogen is led through an acid 


542 Text-Book of Inorganic Chemistry . 

solution. They are : lead, thallium, bismuth, cadmium, tin, and 
copper. 

The metals belonging to the fourth and fifth groups are some¬ 
times called the heavy metals. 

The heaviest or so-called noble metals are : mercury,silver, gold, 
platinum, iridium , palladium, rhodium, ruthenium, and osmium. 
The sulphides of this group are also insoluble in dilute acids. 

We now proceed to the discussion of the metals arranged in 
these six classes. 


The Metals of the Alkalies. 

At the commencement of this century it was universally believed 
that every metal must possess a comparatively high specific gravity 
— i.e. be considerably heavier than water (see p. 331 ). But when 
the metals of the alkalies were discovered by Davy this opinion 
was shown to be incorrect. Potassium and sodium are lighter than 
water, and lithium, which was discovered later, is only about one- 
half as heavy. 

The metals of the alkalies are characterized by their strong 
chemical affinities for negative elements, and are therefore well 
adapted for withdrawing oxygen, sulphur, chlorine, &c., from com¬ 
bination with other elements. They belong to our most powerful 
reducing agents. Water is decomposed by them even at the ordi¬ 
nary temperature of the air. Their oxides and sulphides are easily 
soluble in water and are powerful bases. With the halogens they 
produce, like most other metals, neutral compounds, all of which 
are soluble in water. 


POTASSIUM. 

Chemical Symbol: K .—Atomic Weight : 39 . 

The strong attraction of the metals of the alkalies, and especially 
of potassium, for oxygen, makes it impossible for them to occur in 
the free state in nature. Potassium is chiefly found in the mineral 
kingdom in combination with oxygen and silica as a silicate, and 
in this form is a constituent of potash felspar ; it also occurs, 
united with chlorine, as potassium chloride. The upper layer of 
the salt deposits at Stassfurth in Germany consists principally of 





Potassium. 343 

potassium chloride, which is largely used for the manufacture of 
other potassium compounds. 

Potash felspar (orthoclase), which may be considered as a 
double compound of aluminium silicate and potassium silicate, and 
which is one of the constituents of granite, gradually becomes 
disintegrated under the influence of frost, and is then decomposed 
by water containing carbonic acid into insoluble aluminium silicate 
(clay) and soluble potassium silicate, or carbonate, which fertilizes 
the soil. 

Nearly all land plants require potassium compounds for their 
growth and even for their existence, just as we require salt in our 
food. These potassium compounds are absorbed by their roots 
and converted in their structure into potassium salts of organic 
acids. And when the plants are burnt nearly the whole of the 
potassium which they contain remains behind as potassium car¬ 
bonate or potash. This was formerly the material from which 
potassium compounds were almost exclusively prepared. 

Potassium is a silver-white metal with a bright lustre. Its 
specific gravity is 0 * 87 , it melts at 62 °, and can be distilled at a 
red heat in a stream of hydrogen, giving a vapour of a bright green 
colour. At the ordinary temperature the metal is soft and can be 
easily moulded or cut with a knife, but at o° it is hard and brittle. 
In consequence of its strong attraction for oxygen, the freshly-cut 
surface is only lustrous for a moment, and immediately afterwards 
becomes covered with a thin layer of oxide of a bluish colour. 
The white lustrous surface can only be preserved by melting the 
metal in a glass tube filled with hydrogen, then exhausting the 
gas, sealing up hermetically, and allowing the liquid metal to 
spread over the tube. When slowly cooled, the metal can some¬ 
times be obtained in the crystalline form. 

Potassium is now never prepared by the process by which 
Davy first obtained it—the decomposition of fused caustic potash 
by an electric current—as this only yields a very small quantity 
at one time. It is far better to reduce potassium carbonate with 
charcoal at a red-heat. The chemist is no longer compelled to 
prepare his own potassium, as it is now manufactured in chemical 
works at a far cheaper rate than it can be prepared in small 
quantities. Nevertheless, potassium is still an expensive sub¬ 
stance, principally because it is only used in small quantities in the 
arts. Its price is about 6s. per ounce. 

To obtain a good yield of potassium from potassium carbonate 


344 Text-Book of Inorganic Chemistry. 

and carbon, it is very necessary that the two substances should be 
mixed together as intimately as possible. And as this can scarcely 
be done by mechanical means, it is better to employ some organic 
salt of potassium rich in carbon, of which purified tartar, acid 
potassium tartrate, is the most suitable. The salt is heated in a 
covered crucible, first gently and then to redness, and the carbona¬ 
ceous residue, consisting of an intimate mixture of potassium car¬ 
bonate and charcoal, is then introduced into an iron retort. The 



retort is afterwards heated to whiteness in a furnace, and the potas¬ 
sium vapour which distils over condensed in a suitable receiver. 

One of the wrought-iron bottles in which mercury is imported 
answers well for the retort. It lies in the furnace in a horizontal 
position, supported on iron bars, as shown in fig. 59 . A short 
wide iron tube, open at both ends, is screwed air-tight into the re¬ 
tort and leads to the upper part of the receiver, which contains a 
little naphtha, and is cooled externally by ice. The potassium then 
condenses in the receiver and collects under the naphtha. 

The operation is not without danger, for the iron tube may easily 
























Potassium . 


345 

become plugged with substances less volatile than the potassium, 
and so produce a violent explosion. Great care must therefore be 
taken to keep this tube free from obstruction. This can be done 
by frequently inserting an iron wire bent into a corkscrew shape 
through a hole in the condenser into the tube and moving it back¬ 
wards and forwards. 1 

The potassium which condenses in the receiver is afterwards 
rectified by distilling it in an earthenware retort, or in one of hard 
glass, covered externally with a layer of clay. It is thus obtained 
in small pellets of about the size of a hazel-nut, which, if kept under 
naphtha in well-closed bottles, can be preserved for a long time. 
They are generally coated with a thin brownish crust. 

If potassium is exposed to the air it takes up oxygen, water, and 
carbonic acid, with evolution of heat, and becomes converted into 
a mixture of potassium hydrate (caustic potash) and potassium car¬ 
bonate, which gradually becomes a syrupy solution of potassium 
carbonate :— 

2K + 2H a O = 2KOH + H 2 

and 

2KOH + C 0 2 = CO(OK) 2 + h 2 o. 

When heated in the air potassium burns with a violet flame 
and produces its peroxide. The metal decomposes water at the 
ordinary temperature with a large evolution of heat, and appears 
to burn with a violet flame when thrown upon water. What really 
burns is the hydrogen which the potassium liberates from the 
water in its conversion into potassium hydrate, and which is ignited 
by the large amount of heat evolved during the reaction. The 
violet colour of the burning hydrogen is due to small particles of 
potassium with which it is mixed. 

Potassium also burns in chlorine, and forms potassium chloride. 
At the ordinary temperature it does not unite with hydrogen. But 
at 200°, and especially between 300° and 400°, it begins to absorb the 
gas and becomes changed into a lustrous, brittle, crystalline sub¬ 
stance, which catches fire on exposure to the air, and is again de¬ 
composed at a higher temperature. This compound may have the 
composition : KH. 

1 Violent explosions are also sometimes produced in the manufacture of 
potassium by the formation of a peculiar black compound of potassium and 
carbon which very easily decomposes. These explosions may be prevented 
if the potassium vapour is rapidly cooled by attaching a flattened receiver to the 
iron delivery tube. This exposes a large surface to the air and so rapidly cools 
the hot vapour.—E d. 


346 


Text-Book of Inorganic Chemistry. 


COMPOUNDS OF POTASSIUM. 

Potassium unites with oxygen in two proportions, to form 
potassium oxide : K 2 0 , and potassium peroxide : K 0 2 . In the 
former of these, and all its other compounds except the peroxide, 
the metal plays the part of a monad; in the latter it appears to 
have a higher atomicity. Both compounds are only of theoretical 
interest. 

If potassium is heated in dry oxygen or air, it forms not the 
oxide but the peroxide as a yellow powder, which melts at a higher 
temperature and solidifies on cooling to a crystalline mass. This 
peroxide is decomposed by water, or even by exposure to moist air, 
forming caustic potash, and giving off oxygen. 

Potassium oxide : K 2 0 , which cannot be obtained by heating 
potassium in oxygen^s produced when caustic potash is fused with 
as much potassium as it already contains : KOH + K = K 2 0 + H. 
It is a white powder, which melts at a red-heat, energetically 
attracts moisture from the air, and in contact with water becomes 
converted into potassium hydrate, with a large evolution of heat: 
K 2 0 + H 2 0 = 2 K 0 H. 


Potassium Hydrate, Caustic Potash: KOH. — This important 
compound is a white crystalline solid, melting to a clear liquid at 
a red-heat. It is strongly alkaline and caustic, deliquesces in the 
air, and is easily soluble both in water and alcohol. 

An aqueous solution of pure potassium hydrate may be obtained 
by dropping small pieces of potassium into water. And from this 
solution the solid may be prepared by rapidly evaporating the 
liquid in a silver basin, silver being much less attacked by caustic 
potash than platinum. 

The ordinary method for the preparation of caustic potash on 
the large scale consists in dissolving potassium carbonate in about 
eight or ten times as much water, heating the solution in an iron 
vessel to boiling, and then gradually adding slaked-lime which has 
been previously mixed with water to a thin paste. The lime 
then combines with the carbonic acid, forming calcium carbo¬ 
nate, which is insoluble in water, and caustic potash remains in 
solution. A more concentrated solution of potassium carbonate 
cannot be employed, because it would only be incompletely de- 


Compounds of Potassium. 347 

composed by the lime. Strong caustic potash decomposes calcium 
carbonate when boiled with it, forming potassium carbonate and 
lime. The quantity of quick-lime ' (calcium oxide) which is re¬ 
quired to convert a given weight of potassium carbonate into 
potassium hydrate is calculated according to the following equa¬ 
tion :— 

CO(OK) 2 + CaO + H 2 0 = C 0 - 0 2 Ca + 2KOH, 

and weighed out. This lime is next slaked, mixed with water, 
and the milk of lime dropped into the boiling potassium carbonate 
until a small quantity of the clear liquid no longer effervesces with 
hydrochloric acid—a proof that it no longer contains potassium 
carbonate. The fire is then removed, the vessel covered with an 
iron lid and allowed to cool. As soon as the calcium carbonate 
has settled, and the liquid become clear, it is run off with a 
syphon, and rapidly evaporated in a clean iron or silver dish, when 
solid crystalline potassium hydrate remains behind. 

Potassium hydrate appears to form several crystalline com¬ 
pounds with water, and the solid commercial hydrate usually 
contains more or less water, up to 20 per cent. This water may 
be removed by heating the compound in a silver dish until it fuses 
quietly. A saturated aqueous solution of potassium hydrate de¬ 
posits crystals on cooling, consisting of one molecule of the 
hydrate united with two of water : KOH + 2H 2 0. 

Caustic potash prepared in this way is never pure. It contains 
firstly the impurities present in the commercial potassium car¬ 
bonate, such as potassium chloride, sulphate, &c. ; secondly, the 
impurities of the crude lime ; and, thirdly, potassium carbonate, 
due to the absorption of carbonic acid from the air during the 
evaporation of the solution. To prepare the pure substance from 
the crude hydrate, it is broken into pieces and shaken up in a 
closed vessel with cold alcohol. This liquid dissolves the potas¬ 
sium hydrate, but leaves all the impurities, including the potassium 
carbonate, behind undissolved. The clear solution is then run off, 
diluted with water, and heated, first to drive off the alcohol and 
then more strongly to expel the water. If the alcoholic solution 
were heated alone without the addition of water, the liquid would 
be decomposed under the influence of the potash and the oxygen 
of the air, and would acquire a brown colour and unpleasant 
odour. 

Small quantities of chemically pure potassium hydrate may be 


348 Text-Book of Inorganic Chemistry. 

prepared by boiling a solution of pure potassium carbonate with 
pure lime obtained from Iceland spar or white marble. But as 
caustic potash, as well as its solution, readily dissolves both 
silica and alumina, it can neither be prepared nor evaporated in 
vessels of glass or porcelain; it energetically attacks both these 
substances. 

The fused caustic potash of commerce is the hydrate which 
has been melted and strongly heated as long as it loses water. 
It is sometimes brought into trade in the form of rough flat 
pieces, showing a crystalline fracture, or else as sticks of about the 
thickness of a black-lead pencil, which are obtained by casting 
the fused substance in iron moulds made in two pieces (see 
fig. 60). 



Fig. 60. 


Solid caustic potash rapidly deliquesces in the air, absorbing 
water and carbonic acid, and producing a colourless liquid with a 
strong alkaline reaction. This liquid gives off its carbonic acid 
with effervescence when treated with an acid. 

Aqueous caustic potash gives an oily, slippery feeling to the 
fingers, and possesses a sharp alkaline taste. It powerfully 
attacks the epidermis, and is therefore largely used as a strong 
caustic. The solution dissolves grease and fats, with which the 
alkali unites and forms soluble compounds called soaps ; a dilute 
potash lye (or solution of the alkali) is therefore employed for 
cleansing purposes. 

Caustic potash, both in the solid form and in solution, is 
largely used in the laboratory. In its attraction for acids it 






















Caustic Potash. 


349 


excels all other bases, 1 2 and is, therefore, the most suitable sub¬ 
stance to separate them from their salts. If a solution of a salt of 
copper, for example, is mixed with caustic potash, the correspond¬ 
ing potassium salt is produced, and copper hydrate separates as a 
precipitate :— 

S 0 2 - 0 2 Cu + 2KOH = S 0 2 ( 0 K) 2 + Cu(OH) 2 . 

Caustic potash is also largely used in the arts, but is gradually 
becoming displaced by the cheaper caustic soda, except when the 
latter is too feeble to effect the required change—for example, in 
the conversion of cellulose into oxalic acid (see p. 298).* 


The salts of potassium can be easily prepared by the union of 
the base with the corresponding acid. In this small text-book we 
can only refer to the most important of this numerous class of 
compounds. 

Potassium Sulphate ! S0 2 |qjt is easily obtained by neutral¬ 
izing potassium carbonate with sulphuric acid, and crystallizes 
when a concentrated solution is allowed to cool in hard four-sided 
pyramids or prisms, without water of crystallization. It possesses 
a saline, bitter taste, and melts when strongly heated without 
change. It is not so easily soluble in water as the other potassium 
salts : 100 parts of water at 12 0 only dissolve twelve parts of the 
salt, and at ioo° only twenty-six parts. Alcohol does not dissolve 
it in the least. 


Acid Potassium Sulphate-. S 0 2 |q^ is prepared by heating 

the normal sulphate with the requisite quantity of sulphuric acid, 
and is a bye-product when nitric acid is manufactured from 
potassium nitrate (p. 190). It is easily soluble in water, and 
crystallizes from the hot concentrated solution in rhombohedra; 
it reacts and tastes acid. When heated to about 200 , it melts 
quietly; at a higher temperature it gives off water, and is converted 
into potassium disulphate like the sodium salt (p. 167): 


2 S 0 2 


OH 

OK 


^ (S 0 2 0 K 

u |so;ok 


H 2 0 . 


1 Except the oxides of the rare metals rubidium and caesium. Ed. 

2 Owing to the large quantities of caustic potash now sent into trade from 
the German manufactories, its present retail price scarcely exceeds that of 
caustic soda.—E d. 


350 Text-Book of Inorganic Chemistry. 

At a still higher temperature this salt splits up into sulphuric 
anhydride and the normal sulphate. Hence the use of acid 
potassium sulphate to decompose certain minerals which are not 
attacked by acids. 

Potassium Nitrate ( Nitre , Saltpetre)'. N 0 2 -OK, crystallizes 
from a hot solution in long, furrowed, rhombic prisms, with a 
saline cool taste, and may be easily prepared by neutralizing 
nitric acid with potassium carbonate and evaporating down. 
Nitre is also produced naturally in countries where the soil is 
rich in potassium salts and where rain seldom falls, or only at 
periodic intervals (Egypt, Persia, India, &c.). In these countries 
the surface of the ground after the rains gradually becomes 
covered with a white efflorescence, which consists chiefly of 
potassium nitrate. The salt is then obtained by simple solution in 
water and re-crystallization. This nitre owes its production to the 
potassium salts and nitrogenous organic matter present in the 
soil. The latter gradually undergoes decay, and its nitrogen is 
converted into ammonia, which, in the presence of the strong base, 
becomes oxidized to nitric acid, forming potassium nitrate. 

This process has been imitated in some European countries, 
and large quantities of nitre have been artificially produced in the 
following manner. Heaps of earth mixed with some potassium 
carbonate and lime were made under open sheds and then re¬ 
peatedly watered with liquid manure or other nitrogenous organic 
refuse. The mass was frequently turned so as to expose it as 
much as possible to the oxidizing influence of the air. After a 
considerable period (often one or more years), when all the nitro¬ 
gen had been converted into nitrates, the heaps were lixiviated 
with water to dissolve these soluble salts, and the solution mixed 
with potassium carbonate to convert the calcium and magnesium 
nitrates into the soluble potassium salt. The clear solution was 
then run off from the insoluble calcium and magnesium carbonates 
and evaporated down to crystallize. 1 

Considerable quantities of potassium nitrate are now prepared 
from the sodium nitrate found so largely in Chili. If a solution 

1 This production of nitre from nitrogenous organic substances is exceed¬ 
ingly costly and wasteful, since the value of the nitrogen in the form of am¬ 
monia for increasing the fertility of the soil is far greater than that of the nitre 
obtained from it, and the process has now been almost entirely abandoned.— 
Ed. 


Potassium Nitrate. 


35i 


of this salt is mixed with one of potassium chloride, and the 
mixed solution evaporated down under pressure, double decom¬ 
position occurs into sodium chloride (common salt) and potassium 
nitrate. And since sodium chloride is only slightly more soluble 
in hot than in cold water, this salt first separates out on evaporating, 
while potassium nitrate, which is much more soluble in hot than 
in cold water, crystallizes out on cooling the hot decanted liquid. 

Potassium nitrate prepared in any of these ways is always 
mixed with impurities, especially common salt, from which it can 
be freed by repeated crystallization. This purification depends 
upon the greater solubility of saltpetre in hot than in cold water. 
100 parts of water dissolve only 13 parts of the salt at o°, and 29 
parts at 18 0 , but at 97 0 the same quantity of water dissolves 236 
parts. And a boiling-hot saturated solution of saltpetre, of which 
the boiling-point is 116 0 , is said to contain 335 parts of the salt 
dissolved in 100 parts of water. 

Nitre is used for various purposes in medicine, and sometimes 
for the manufacture of nitric acid, but far larger quantities are ab¬ 
sorbed in the manufacture of gunpowder, in which the nitre is 
the oxidizing substance. The well-known explosive action of gun¬ 
powder is due to the sudden production of large volumes of hot 
gases from a small volume of a solid. If powdered nitre is 
sprinkled on a piece of glowing charcoal, the latter burns brilliantly 
to carbonic acid and carbonic oxide, while the nitre is converted 
into potassium carbonate and free nitrogen. The same process 
goes on when an intimate mixture of powdered nitre and charcoal, 
in the correct proportions, is ignited, and may be represented by 
the following equation :— 

2N0 2 -0K + 3C = CO, + CO + N 2 + 

N Gunpowder contains, however, sulphur as well as nitre and 
charcoal, and it was formerly thought that its decomposition when 
burnt was expressed by the equation 

2NCVOK + 3C + S = 3 C 0 2 + N 2 + K 2 S, 
which represents the various constituents in the proportions in 
which they are actually contained in gunpowder. But it has been 
found that other substances, besides those shown in the above 
equation, are also produced, and that the sulphide formed is 
potassium disulphide and not the monosulphide. The substances 
into which gunpowder is converted when burnt in a closed space 
are essentially:— 



352 Text-Book of Inorganic Chemistry. 

(i.) Solids , includingpotassium carbonate, potassium sulphate, and 
potassium disulphide, and which are about 57 per cent, of the whole. 

(ii.) Gases which consist of carbonic acid, nitrogen, and carbonic 
oxide, and constitute the remaining 43 per cent, by weight. The 
volume of these gases, measured at the normal temperature and 
pressure, is nearly 300 times that of the powder, and at the high 
temperature of the explosion nearly 2,000 times. Hence the force 
of the explosion. Other substances are also formed when gun¬ 
powder is exploded, but only in very small quantities. And the 
actual chemical changes which go on when the powder is burnt 
can only be represented by a somewhat complex equation. 

The gunpowder of different countries, and even of the same 
country, varies considerably. The mean composition of English 
powder is about as follows :— 

Nitre.753 

Charcoal . . . . . .13-4 

Sulphur.n*3 

1 oo-o 

In the manufacture of a suitable gunpowder it is not only 
necessary to mix the constituents intimately together, but the 
mixture must be afterwards granulated, an operation which is 
performed by suitable machinery. The size of the grain depends 
upon the purpose far which the powder is to be used : it varies 
from abour tnc size of a millet seed for ordinary purposes and 
small arms, up to blocks as large as a cubic inch (pebble powder) 
for large cannon. 

By diminishing the quantity of nitre, and increasing that of 
charcoal or sulphur, the rate at which the powder burns is dimin¬ 
ished. And as the gases produced by the burning of the powder 
cannot support combustion, but, on the contrary, extinguish bodies 
in the act of burning, this property has been successfully utilized 
to extinguish fires in closed places. A mixture for this purpose 
may contain, for example, 60 to 66 per cent, nitre, 30 to 36 per 
cent, sulphur, and at most 4 per cent, charcoal. On burning, it 
produces sulphurous anhydride with carbonic acid and nitrogen, 
all of which extinguish burning bodies. 

Potassium Nitrite'. NO-OK.—If potassium nitrate is melted at 
a low red heat, it is decomposed into oxygen which escapes, and 
potassium nitrite which remains behind, usually mixed either with 


Potassium Carbonate. 


353 

undecomposed nitrate, or, if too strongly heated, with caustic 
potash. The salt is easily soluble in water, deliquesces in the air, 
and evolves copious red vapours of nitrous acid when treated with 
moderately dilute sulphuric acid. 


Potassium Carbonate : CO | q deliquesces in the air and 

is easily soluble in water. The aqueous solution reacts strongly 
alkaline, and when evaporated down leaves the salt as a white solid 
which melts unchanged at high temperatures. An extremely con¬ 
centrated, warm solution deposits crystals containing about 16 per 

cent, of water of the composition: 2CO | + 3H 2 0. 

Potassium carbonate is a constituent of the ashes which remain 
when wood and land plants generally are burnt, which, however, 
do not contain more than 20 per cent, of it. The ashes consist 
chiefly of the sulphates, phosphates, and silicates of potash, lime, 
and magnesia, together with some chlorides. 

Crude potash is obtained by extracting the ashes with water, 
and evaporating the filtered solution to dryness. It contains at 
most 60 per cent, of potassium carbonate mixed with the other 
soluble salts contained in the ashes. By digesting this crude 
potash with an equal weight of water for several days, with frequent 
stirring, the easily soluble carbonate is dissolved, leaving most of 
the other salts behind. The clear solution evaporated to dryness 
gives the pearlask of commerce. 

This is still very far from pure potassium carbonate, and in fact 
we know of no method of completely purifying it. Chemically 
pure potassium carbonate may, however, be easily obtained by 
heating a pure potassium salt of an organic acid in a platinum 
vessel. Vessels of porcelain or glass cannot be employed, as the 
hot potassium carbonate abstracts silica from them, and so becomes 
impure. For the preparation of the pure salt, acid potassium 
tartrate is commonly used. This salt can be readily obtained, and, 
as it is only difficultly soluble in water, can be easily purified. 
The heated mass, when extracted with water, and the clear solution 
evaporated in a platinum dish, leaves pure potassium carbonate as 
a snow-white solid. 

Since the discovery of the potassium chloride beds at Stassfurth, 
potassium carbonate is manufactured from the chloride, just as soda 
is obtained from common salt by Leblanc’s process (seep. 371). 

* A A 



354 Text-Book of Inorganic Chemistry. 

Potassium carbonate is one of the most important of the 
potassium salts. It is used not only for the preparation of caustic 
potash and many other potassium compounds, but is also largely 
employed in the arts, especially in the manufacture of glass and of 
soft soap. 

( OPT 

Acid Potassium Carbonate'. CO j 0K is obtained by leading 

carbonic acid into a concentrated aqueous solution of the normal 
salt, when it separates out as a crystalline powder:— 

CO {oK + C ° 2 + Hs ° = 3 «>{8i 

It requires about four times its weight of water for solution, reacts 
neutral, and is decomposed at about 8o° into the normal salt and 
carbonic acid. 

Potassium Chlorate: C 10 2 - 0 K.—This salt, of which the pre¬ 
paration has been already described (p. iio), crystallizes when its 
hot saturated solution is cooled in lustrous scales or larger plates. 
It melts at 334 0 , and begins to be decomposed, with evolution of 
oxygen, at 352 0 . It requires 16 parts of water for solution at the 
ordinary temperature, but only 2 parts at ioo°. In consequence of 
the readiness with which it parts with oxygen, it is an excellent 
oxidizing agent, and yields considerable quantities of chlorine when 
warmed with hydrochloric acid (p. 112). 

Potassium Perchlorate : C 10 3 - 0 K, has been already described. 
It belongs to the potassium salts which are difficultly soluble in 
water, and can, therefore, be readily purified by crystallization. 
The salt is insoluble in alcohol. 

Potassium Hypochlorite : CIOK.—This salt, which is only 
known in its aqueous solution and which is characterized by its 
powerful bleaching properties, is obtained by passing chlorine 
into a cold, dilute solution of potassium carbonate until large 
quantities of carbonic acid are evolved. Potassium chloride and 
hypochlorite with acid potassium carbonate are first produced 

2C °{OK + 2 C 1 + H 2 0 = 2C0|0K + KC 1 + CIOK, 

and carbonic acid is, therefore, not evolved at once, but only when 
the chlorine begins to attack the acid potassium carbonate. 


Potassium Chlorate . 355 

This solution of potassium hypochlorite is known as eau de 
Javelle. Strong acids liberate considerable quantities of chlorine 
from it. 


Potassium Oxalate : {^q-OK + H 2°- The normal salt is ob¬ 
tained by neutralizing oxalic acid, or the acid salt, with potassium 
carbonate, and easily crystallizes from its aqueous solution. The 
(CO • OH 

acid salt : (co-OK + ^2^ {salts of sorrel ), occurs in some 

plants, especially in the leaves of the sorrels, from which it 
may be obtained in large crystals by pressing the leaves and 
evaporating the acid juice. It has a strong acid reaction, and 
unites with ferric oxide to form a soluble double salt. Hence its 
use to remove ink or iron stains from paper or linen. After the 
stains have been removed, the fabric must be carefully washed, 
first with water and then with dilute soda, to prevent corrosion by 
the oxalic acid contained in the acid salt. Besides this acid salt, 
a so-called diacid salt (potassium quadroxalate) is also known! 
The salt crystallizes well, and is to be considered as a double 
compound of acid potassium oxalate and oxalic acid 


j CO • OK 
(CO-OH + 


jCO-OH 

(CO-OH 


+ 2 H, 0 . 


Potassium Silicate (potash water-glass).—Silica and caustic 
potash or potassium carbonate may be fused together in varying 
proportions, and produce, when the quantity of silica is not too 
great, a glass which is distinguished from ordinary glass by its 
solubility in water. Hence the name water-glass. It may be 
obtained by melting together equal weights of p'earlash and sand 
until the mass fuses quietly, or by dissolving infusorial earth in 
caustic potash. By the latter operation a solution is at once 
obtained. In England considerable quantities of water glass are 
manufactured by heating powdered flints with caustic potash (or 
soda) under pressure, when the flints readily dissolve. On 
evaporating a solution of water-glass, an amorphous transparent 
mass remains behind, which, when exposed to the air, gradually 
becomes covered with an opaque coating of silicic acid and 
potassium carbonate. The solid fused water-glass undergoes the 
same decomposition. 

An aqueous solution of water-glass has many technical appli¬ 
cations. It is used, for example, to render theatrical scenes and 


A A 2 


356 Text-Book of Inorganic Chemistry. 

other stage appliances uninflammable. Such articles when coated 
with a thin layer of water-glass are no longer combustible; the thin 
layer of potassium silicate protects their surface like a varnish 
from the oxygen of the air. If brought into contact with a burn¬ 
ing body they become charred but do not catch fire. Water-glass 
is also largely used for protecting buildings from damp, for 
preventing the decay of woodwork in damp situations, and very 
largely in the manufacture of the cheaper kinds of soap. 

In a very dilute aqueous solution potassium silicate forms a 
valuable manure for many plants, especially for those which 
require potash salts and silica for their growth. Pasture lands are 
improved by occasionally watering with a dilute solution of potash 
water-glass. 

It has already been mentioned that potassium silicate is a 
constituent of many minerals—for example, of felspar. 

Potassium Chloride : KC 1 , crystallizes like sodium chloride 
in the regular system, and usually in cubes without water of 
crystallization. The crystals contain, however, water mixed 
mechanically, which causes them to decrepitate when heated. It 
has a saline taste, and is tolerably soluble in water : ioo parts of 
water dissolve about 33 parts of the salt at the ordinary tempera¬ 
ture, and hot water only a little more. It melts without change 
when heated to redness, and then volatilizes in not inconsiderable 
quantities. Potassium chloride is a constituent of sea-water, is 
contained in the ashes of land plants, and is also found in large 
quantities as a mineral. It is the chief constituent of the salt 
beds in the neighbourhood of Stassfurth in Northern Germany. 
The mineral sylvin consists almost exclusively of it. 

Potassium Bromide : KBr, closely resembling potassium 
chloride, may be prepared by neutralizing potassium carbonate 
with hydrobromic acid, or by saturating caustic potash with bro¬ 
mine, evaporating to dryness, and then glowing to destroy the 
potassium bromate produced at the same time. It is now largely 
used in medicine for neuralgic affections. 

Potassium Iodide : KI.—This salt resembles potassium chlo¬ 
ride in its external properties. Like the chloride, it crystallizes in 
cubes without water. It is, however, much more easily soluble in 


Potassium Iodide. 


35 7 

water than potassium chloride, and thus remains in the mother- 
liquor when a mixed solution of the two salts is evaporated down. 
At the ordinary temperature ioo parts of water dissolve more than 
140 parts of potassium iodide ; the saturated solution contains, at 
its boiling-point (120°), more than 200 parts of potassium iodide for 
every 100 parts of water. 

Potassium iodide is of such importance in medicine and photo¬ 
graphy that many attempts have been made to discover profitable 
and ready methods for preparing it. The simplest method—viz. 
neutralization of caustic potash or potassium carbonate with hydri- 
odic acid,just as potassium chloride maybe obtained—is impracti¬ 
cable, because hydriodic acid is very much more difficult to obtain 
in quantity than hydrochloric acid. 

Two methods are chiefly used for the preparation of potassium 
iodide. One consists in gradually adding iodine to warm caustic 
potash until the liquid becomes yellow from free iodine, evaporating 
the liquid, which now contains potassium iodide and iodate (p.130), 
to dryness, and glowing, in order to convert the potassium iodate 
into potassium iodide and oxygen. It has been found, however, that a 
much higher temperature is required to decompose potassium iodate 
than potassium chlorate, and at this high temperature some of the 
potassium iodide is volatilized and lost. It is therefore customary 
to add a little powdered wood charcoal to the liquid containing the 
mixed salts. When the saline mass is afterwards glowed, this 
charcoal abstracts all the oxygen from the iodate and is converted 
into carbonic acid. The residue is then extracted with water, and 
the clear solution evaporated down to crystallize. 

A second method largely used for the preparation of potassium 
iodide is as follows. One part of iron filings is mixed with water 
and digested with three parts of iodine, when the two substances 
readily unite with one another, with considerable rise in temperature, 
and produce the easily soluble salt—ferrous iodide : Fel 2 . As 
soon as all the iodine has disappeared, the green solution is 
filtered, and one more part of iodine added, which produces the 
compound ferroso-ferric iodide, Fel 2 , Fe 2 I 6 . The liquid is then 
exactly precipitated with potassium carbonate, when ferroso-ferric 
hydrate easily separates and can be readily washed, and the clear 
solution, which now contains potassium iodide, is evaporated down 
to crystallize. From 4 parts of iodine, 5 parts of potassium iodide 
may be thus obtained. The ferrous iodide might be at once 
decomposed by potassium carbonate, but the precipitate then 


358 Text-Book of Inorganic Chemistry. 

consists of ferrous carbonate, which does not readily settle, and 
can only be washed with difficulty, and therefore requires large 
quantities of water and a long time to extract all the potassium 
iodide. 

An aqueous solution of potassium iodide dissolves large quan¬ 
tities of free iodine; the saturated solution contains about twice as 
much iodine as was present in the original salt, producing the com¬ 
pound : KI 3 , which may be obtained as dark, lustrous needles, on 
careful evaporation. This compound is very deliquescent, and 
easily decomposes into potassium iodide and iodine. 

Potassium iodide, together with the sodium salt, is present in 
many mineral waters, and especially in the ashes of sea plants 
(pp. 123, et seq.). It is largely used in medicine as a remedy for 
goitre and swellings, in cases of syphilis, &c., and is, in fact, one of 
the most highly prized medicines. Large quantities of the salt are 
also used in photography. 

Potassium Fluoriile : KF, is easily obtained by neutralizing 
hydrofluoric acid with potassium carbonate. It crystallizes at the 
ordinary temperature with two molecules of water, but above 35 0 
without water, in cubes; the salt easily deliquesces in moist air, 
and strongly attacks glass. 

Potassium fluoride unites with other fluorides and with hydro¬ 
fluoric acid, and produces double salts, in which the fluorine appears 
to play the part of a dyad. 

Acid Potassium Fluoride : HKF 2 .—This salt is easily prepared 
by neutralizing a given volume of aqueous hydrofluoric acid with 
potassium carbonate, and then adding an equal volume of the 
acid. It crystallizes out when the solution is evaporated down 
in a platinum basin. The crystals are anhydrous; they easily 
dissolve in water, but difficultly in that containing hydrofluoric 
acid. It requires heating to redness before it is decomposed into 
potassium fluoride and hydrofluoric acid. 

Potassium Fluoborate\ KF,BF 3 = KBF 4 , is precipitated as a 
gelatinous, iridescent mass, which dries to a white impalpable 
powder when potassium fluoride, or any other soluble potassium 
salt, is mixed with aqueous fluoboric acid. It is difficultly soluble 
in cold water, of which it requires 70 times its weight for solu¬ 
tion. 


Potassium Cyanide . 359 

Potassium Fluosilicate : 2KF,SiF 4 = K 2 SiF 6 , closely resembles 
the fluorborate, but is still less soluble in water. Potassium fluo¬ 
silicate is therefore sometimes used as a test for potassium. 

Potassium Cyanide : KCN = KCy.—This extremely poisonous 
salt is decomposed by all acids, including carbonic acid, giving 
free hydrocyanic acid. It may be obtained pure by passing hydro¬ 
cyanic acid gas into well-cooled alcoholic potash, until the mass 
becomes pasty from separation of the potassium cyanide, which is 
almost insoluble in alcohol. The crystalline mass is then brought 
on a filter, quickly washed with strong alcohol, as much of the 
liquid removed as possible by a filter-pump, and finally dried in a 
vacuum over sulphuric acid. 

The less pure compact potassium cyanide of commerce, which 
occurs as flat pieces with a crystalline fracture, is prepared by a 
much less expensive process. Potassium ferrocyanide : K 4 FeCy 6 + 
3H 2 0, which may be considered as a double cyanide of ferrous 
iron and potassium : 4KCy, FeCy 2 + 3 H 2 0 , is first thoroughly freed 
of its water of crystallization by powdering it and heating it in a 
shallow iron dish. The dried salt is then placed in an iron 
crucible, which it must not more than two-thirds fill, and heated 
to fusion in a furnace. The ferrous cyanide is then decomposed 
into nitrogen gas which is given off, and black iron carbide which 
remains behind. The potassium cyanide melts unchanged, pro¬ 
vided it is not exposed to the air, from which (when in the fused 
state) it readily absorbs oxygen, and becomes partly converted 
into potassium cyanate. As soon as the salt fuses quietly, the 
crucible is removed from the furnace, and repeatedly knocked on 
some hard body that the iron carbide may settle to the bottom and 
the potassium cyanide become clear. When the fused salt has 
sufficiently clarified, which may be easily known by removing a 
drop on the end of a glass rod, the liquid is carefully poured on to 
an iron slab, broken into pieces when cold, and preserved in well- 
closed vessels. 

In this method of preparation two of the six atoms of cyanogen 
contained in a molecule of potassium ferrocyanide are lost: the 
two atoms combined with the iron are decomposed. In order to 
obtain this cyanogen as well, it is best to use Liebig’s method, 
which consists in mixing every 8 parts of the dry potassium ferro¬ 
cyanide with 3 parts of potassium carbonate. When this mixture 
is melted, the ferrous cyanide of the potassium ferrocyanide and 


360 Text-Book of Inorganic Chemistry. 

the potassium carbonate first form ferrous carbonate and potassium 
cyanide :— 

FeCy 2 + CojgK = 2 KCy + CO-0 2 Fe. 

But the ferrous carbonate is at once broken up at the high tem¬ 
perature into carbonic acid and ferrous oxide :— 

CO • 0 2 Fe = C0 2 + FeO, 

the oxide being finally reduced to metallic iron by a portion of 
the potassium cyanide, which is converted into potassium cyanate. 
Thus, of the two atoms of cyanogen which were united with the 
iron, one is obtained as potassium cyanide and one as potassium 
cyanate. 

As this small quantity of potassium cyanate is without im¬ 
portance for most of the technical purposes for which potassium 
cyanide is used, the commercial salt is now always manufactured 
by Liebig’s process. The quantity obtained is not only larger, but 
the product is also whiter. 

Potassium cyanide is very easily soluble in water, and deli¬ 
quesces in moist air. It cannot be crystallized from its aqueous 
solution, much less obtained pure, as it gradually decomposes with 
water into potassium formate and ammonia :—. 

KCN + 2H s O = jcO'OK + NH »- 

It is scarcely soluble in alcohol. The aqueous solution reacts 
alkaline, and plays the part of a base with many other insoluble 
cyanides. Just in the same way as potassium hydrate dissolves 
lead oxide, alumina, and silica, producing salts in which these 
substances take the part of an acid, so, too, potassium cyanide 
dissolves lead cyanide, silver cyanide, and numerous other similar 
insoluble compounds, forming salts which are soluble in water, 
and can usually be obtained in the crystalline state. 

Neither solid potassium cyanide nor its aqueous solution unites 
directly with oxygen. But in the fused state its attraction for 
oxygen is so great that most of the metallic oxides are decomposed 
by it. The substance produced by this oxidation is potassium 
cyanate. The salt also easily combines directly with sulphur, 
forming potassium sulphocyanate. Union of the two substances 
even takes place when flowers of sulphur are stirred up with a 
warm aqueous solution of potassium cyanide. 


Potassium Cyanate. 361 

Potassium Cyanate : CONK.—This white crystalline salt is 
easily soluble both in water and alcohol, and may be purified by 
crystallizing from the latter liquid. It is obtained by adding 
an oxide of lead to fused potassium cyanide, when the oxide is 
reduced to metallic lead and the cyanide oxidized to cyanate. It 
is, however, best prepared directly from potassium ferrocyanide by 
oxidization with manganese peroxide. 

Well dried and finely powdered potassium ferrocyanide is 
thoroughly mixed with an equal weight of dry commercial manga¬ 
nese peroxide containing about 75 per cent, of the pure compound, 
and the mixture gently heated with continual stirring in a flat iron 
dish. After a little time the mass begins to glow ; the heating 
is continued as long as this lasts and until a fragment when 
dissolved in water gives no precipitate of Prussian blue on mixing 
the clear solution with ferric chloride. 

The mass is next repeatedly digested wdth strong alcohol (82 
per cent.), and potassium cyanate then separates from the hot 
saturated solution on cooling as thin tablets. After filtering off 
the crystals, they must be rapidly dried, as potassium cyanate 
readily decomposes with water, even at the ordinary temperature, 
into potassium and ammonium carbonates :— 

2CONK + 4H 2 0 = CO(OK) 2 + CO(ONH 4 ) 2 . 

Its aqueous solution, when mixed with ammonium sulphate, 
forms potassium sulphate and ammonium cyanate, which latter 
changes immediately into the isomeric compound urea. Addition 
of a strong acid to a solution of potassium cyanate does not 
liberate free cyanic acid, but decomposes it into carbonic acid and 
ammonia (p. 317). 

Potassium Suiphocyanate\ CSNK, may be obtained by adding 
sulphur to fused potassium cyanide, or, better, by heating a mix¬ 
ture of 46 parts of potassium ferrocyanide, 17 parts of potassium 
carbonate, and 32 parts of sulphur in a Hessian crucible until 
the mass fuses quietly, and until a drop dissolved in water no 
longer gives a blue precipitate with ferric chloride, but only a red 
colour. The fused mass is broken in pieces when cold and 
digested with hot alcohol, which dissolves the potassium sulpho- 
cyanate and leaves ferrous sulphide behind. From the hot 
solution it crystallizes on cooling in long colourless crystals, 
resembling nitre, and possessing, like this salt, a cooling taste. 


362 Text-Book of Inorganic Chemistry. 

Water dissolves large quantities of it, so that it deliquesces in 
moist air. A considerable fall in temperature is produced when 
the salt is dissolved in water. 

Potassium sulphocyanate, or any soluble sulphocyanate, gives 
no precipitate with ferric salts, but a deep red colour. 


Sulphur Compounds of Potassium. 


The alkali metals unite with sulphur in several proportions, 
forming definite chemical compounds, which may be compared 
with the compound of potassium iodide with iodine. Just as potas¬ 
sium iodide unites with two atoms of iodine (p. 359), so also potas¬ 
sium sulphide unites with one, two, three, and four atoms of sulphur, 
producing the compounds :— 


Potassium disulphide 
Potassium trisulphide 
Potassium tetrasulphide . 
Potassium pentasulphide 


' -^ 2^2 
. k 2 S 3 
. k 2 S 4 

. k 2 s 5 . 


The two atoms of potassium do not appear to combine with more 
than five atoms of sulphur. 


Potassium Monosulphide : K 2 S, is a strong sulpho-base, easily 
soluble in water, and deliquescing in the air. It has a strong 
alkaline reaction and taste. This compound may be obtained by 
reducing potassium sulphate in a stream of hydrogen at a red- 
heat :— 

S 0 2 (OK) 2 + 4 H 3 = K 2 S + 4 H 2 0 , 

or, better, by glowing a mixture of one molecule of the salt with four 
molecules of carbon :— 

S 0 2 ( 0 K) 2 + 4C = K 2 S + 4CO. 

A mixture of 3 parts of powdered potassium sulphate and 1 part 
of wood charcoal is placed in a covered Hessian crucible, which it 
must not more than two-thirds fill, and heated until the mixture has 
ceased frothing and fuses quietly. When the crucible is cold it is 
broken, and the potassium sulphide obtained as a red crystalline 
mass. It soon loses its colour in the air, from which it abstracts 
not only water, but also oxygen and carbonic acid. 

An aqueous solution of potassium monosulphide is readily ob 
tained in the following way. A given volume of moderately con¬ 
centrated caustic potash is divided into two equal parts. One of 


Sulphur Compounds of Potassium. 363 

these is then completely saturated with sulphuretted hydrogen gas, 
which forms the sulphydrate: KSH, and the other part is then 
added. One molecule of potassium sulphydrate is decomposed by 
one molecule of caustic potash into one molecule of potassium 
monosulphide and one of water - 

KSH + KOH = K 2 S + H 2 0 . 

This solution is decomposed when treated with dilute acids, 
and then evolves sulphuretted hydrogen without depositing sulphur. 
It absorbs oxygen from the air, and is converted into potassium 
thiosulphate :— 

K 2 S + 30 = SO, | Of 

Potassium monosulphide easily dissolves the sulphides of 
arsenic, antimony, and carbon and other similar sulpho-acids, with 
which it unites chemically and forms soluble sulpho-salts. 

Potassium Sulphydrate : KSH or ^ j S. This compound, which 

corresponds to caustic potash, is obtained by leading sulphuretted 
hydrogen into caustic potash as long as the gas is absorbed. If the 
solution of caustic potash is sufficiently concentrated, the compound 
separates out on cooling in long colourless crystals, containing 
water of crystallization. In its chemical behaviour it closely re¬ 
sembles the monosulphide. 

The poly sulphides of potassium are obtained by adding the 
requisite quantity of sulphur to the fused monosulphide, or to its 
concentrated aqueous solution. If more sulphur is added to fused 
potassium monosulphide than corresponds to 4 molecules of the 
former to 1 molecule of the latter, there is produced 

Potassium Pentasulphide : K 2 S 5 , and the excess of sulphur 
volatilizes. This compound has a yellow-brown colour; it easily 
dissolves in water, and when acted on by hydrochloric acid yields 
sulphuretted hydrogen and large quantities of milk of sulphur 
(p. 145). When heated more strongly it loses sulphur and be¬ 
comes 

Potassium Trisulphide : K 2 S 3 , a compound closely resembling 
the pentasulphide. 

liver of Sulphur ( Potassa sulphuratal) —This pharmaceu¬ 
tical preparation, which owes its name to its colour, consists essen- 


364 Text-Book of Inorganic Chemistry . 

tially of a mixture of potassium pentasulphide and potassium 
sulphate. It is prepared by gradually heating a mixture of about 
equal parts of potassium carbonate and sulphur in a large earthen¬ 
ware crucible. At first considerable quantities of carbonic acid 
are evolved, which causes the mixture to froth up ; as soon as this 
has ceased, the crucible is closed and the temperature gradually 
raised until all fuses quietly. The liquid is then poured on to a 
sheet of iron, broken into pieces when cold, and quickly brought 
into well-closed bottles, as it readily attracts moisture from the air. 
The following equation represents the reaction :— 

4CO(OK) 2 + 16S = 3K 2 S 5 4 S 0 2 ( 0 K) 2 + 4 C 0 2 . 

The preparation, when well made, should dissolve in water to a 
nearly clear solution, and when mixed with hydrochloric acid should 
give sulphuretted hydrogen and large quantities of free sulphur. 

A solution containing potassium pentasulphide and thiosul¬ 
phate may be obtained by gradually adding flowers of sulphur to 
boiling caustic potash as long as the sulphur is dissolved : — 

6KOH + 12S = 2K 2 S 5 + S0 2 |OK + 3 h 2 0 . 


Detection of Potassium Compounds. 

Potassium compounds when heated in the Bunsen burner 
impart a characteristic violet colour to the flame. In the presence 
of sodium, even in small quantities, the intense yellow light emitted 
by this substance obscures the feeble potassium flame, and it is 
therefore not seen by the naked eye. But if the flame is viewed 
through a piece of blue cobalt glass, the yellow sodium light is 
absorbed by the blue of the glass, while the violet potassium light 
passes through the glass unchanged. The most certain way of 
detecting traces of potassium is by its spectrum, which consists 
essentially of two lines in the red and one in the violet. (See Table 
of Spectra). 

If a solution of potassium chloride or of any other soluble 
potassium compound (except the iodide) is mixed with platinic 
chloride, a yellow crystalline precipitate of a double chloride of 
potassium and platinum : 2KCI, PtCl 4 , or potassium chlorplatinate : 
K 2 PtCl 6 , is produced, which is slightly soluble in water, but in¬ 
soluble in alcohol. The heavy metals, as well as the metals of 
the earths and alkaline earths, if present, must have been pre- 


Detection of Potassium Compounds. 365 

viously removed from the solution, and ammonium compounds 
destroyed by glowing. Sodium salts do not give this precipitate, 
as the corresponding sodium salt is soluble both in water and 
alcohol. This double chloride of potassium and platinum is 
decomposed on glowing into metallic platinum, gaseous chlorine, 
and potassium chloride. The last-named compound can be ex¬ 
tracted with water, and easily obtained in cubic crystals. 

If a concentrated solution of a potassium salt is mixed with a 
concentrated solution of tartaric acid in excess, and the mixture 
well shaken or stirred with a glass rod, a white crystalline pre¬ 
cipitate of acid potassium tartrate is produced, which is first 
deposited on those parts of the glass vessel which have been 
touched by the rod. Sodium salts do not produce this precipitate, 
as the corresponding acid sodium tartrate is easily soluble in 
water. 


SODIUM. 

Chemical Symbol : Na.— Atomic Weight: 23. 

Sodium, which is. closely allied to potassium, is also found in 
nature in the form of a double silicate of aluminium as soda felspar 
or albite. It further occurs as the mineral cryolite : 6NaF, A 1 2 F 6 , 
which is worked both for soda and for alumina. But by far its 
most important natural compound is the chloride —common salt or 
rock-salt. 

The preparation of sodium is carried on in just the same way 
as that of potassium, except that some powdered chalk is usually 
added to the mixed sodium carbonate and charcoal, to prevent 
fusion at the high temperature necessary for reduction. Since the 
affinity of sodium for oxygen is less than that of potassium, its 
reduction, according to the method described on p. 344, is easier 
and the yield is greater. There is also less danger of the formation 
of explosive compounds, as in the preparation of potassium. This 
is one reason why sodium is so much cheaper than potassium. 

It is chiefly owing to the researches of St. Claire Deville, sup¬ 
ported by Napoleon III., who, in attempting to prepare cheap 
aluminium by means of sodium, so perfected the method for the 
extraction of this metal that its price has fallen from $s. per oz. to 
Ss. per lb. during the last thirty-five years. The present low price 



356 Text-Book of Inorganic Chemistry. 

of sodium is also partly due to the large quantities which are now 
used for the manufacture of metallic magnesium. 

Commercial sodium comes into trade in the form of bars. 
When freshly cut the surface is of a silver-white colour, but it very 
rapidly tarnishes in the air, becoming covered with a layer of its 
oxide. At the ordinary temperature it is soft like wax, and can be 
easily moulded or pressed into any shape. The metal is lighter 
than water, but somewhat heavier than potassium—its specific 
gravity being 0-978. Its melting-point, which is also higher than 
that of potassium, is 96°. It catches fire when heated in the air, 
and burns with an intensely yellow flame. Water is easily decom¬ 
posed by it with evolution of hydrogen, and the sodium hydrate 
which is produced at the same time dissolves in the water and gives 
it an alkaline reaction. The liberated hydrogen does not, however, 
ignite like that set free by the action of potassium on water, because 
the quantity of heat set free when sodium decomposes water is less 
than when the same decomposition is effected by potassium—the 
affinity of the former metal for oxygen being less than that of the 
latter. The liberated heat would, however, raise the temperature 
of the hydrogen to its ignition point, if it were not continually 
cooled by the movement of the fragment of sodium on the surface 
of the cold water. If hot water, at about 8o°, instead of cold water, 
is used for the experiment, the hydrogen catches fire at once and 
burns with a yellow flame. Or, if the surface of the water is covered 
with a piece of filter-paper the movement of the pellet of sodium 
is prevented, and it remains in contact with only a small quantity 
of water ; this then becomes so strongly heated that the hydrogen 
catches fire. 

Potassium and sodium unite together and form an alloy which 
is liquid at the temperature of the air, and closely resembles 
mercury. This is a striking example of the law that the melting- 
point of an alloy is lower than that of its constituents. 

Sodium, like potassium, unites chemically with hydrogen when 
heated up to about 300° in the gas. This compound is again de¬ 
composed at about 420°. If dry ammonia is led over gently heated 
sodium the gas is decomposed with the formation of sodium amide : 
Na-NH 2 , while hydrogen is set free. When fused, this compound 
has a greenish-blue colour, but solidifies to a pink crystalline 
mass. 



Compounds of Sodium . 


367 


COMPOUNDS OF SODIUM. 

With few exceptions, the compounds of sodium closely resemble 
those of potassium, and we shall, therefore, only describe those in 
detail which differ essentially from the corresponding potassium 
compounds. 

Sodium, when heated in dry oxygen or air, produces sodium 
peroxide : ^ Na 2 0 2 , similar in its properties to the potassium com¬ 
pound. Sodium oxide : Na 2 0 , is obtained, like potassium oxide, 
by heating sodium hydrate with sodium. 

Sodium Hydrate (Caustic Soda ): NaOH.—This compound is 
so similar to potassium hydrate that the two cannot be externally 
distinguished from one another. The general behaviour and the 
preparation of the two substances is also the same, except that 
caustic soda is a less powerful base than caustic potash. In pre¬ 
paring caustic soda from crystalline sodium carbonate, according 
to the method described on p. 346, it must be remembered that the 
sodium salt contains about 60 percent, of water, and not more than 
five parts of water should, therefore, be used to dissolve one part 
of the salt. 

Sodium Sulphate: S 0 2 |o>Ja + IoH 2 0 -—This salt, com¬ 
monly known as Glauber's salt , is occasionally found in the solid 
state in nature, but usually in solution in various mineral springs. 
Many of these springs, especially those of Marienbad, Carlsbad, 
and others, are highly valued on account of the mild aperient 
action of the sodium sulphate which they contain. Sodium sul¬ 
phate may be easily obtained by neutralizing sodium carbonate 
with sulphuric acid, or by heating this acid with common salt. 
This latter process is largely carried on in alkali (soda) works, 
where common salt is first converted into sodium sulphate prior to 
its change into sodium carbonate. 

Sodium sulphate crystallizes with ten molecules (i.e. 56 per cent, 
of water) in large colourless prisms, which, when heated, melt in 
their water of crystallization, and finally lose all their water. It is 
easily soluble in water but insoluble in alcohol. At o°, 100 parts 
of water dissolve 12 parts of the salt, at 18 0 48 parts, at 25 0 100 


368 Text-Book of Inorganic Chemistry . 

parts, and at 33 0 as much as 322 parts. But above 33 0 the solu¬ 
bility diminishes, because the compound with ten molecules of 
water cannot exist above this temperature. It is then converted 
into the anhydrous salt, which is less soluble in water, and separates 
from the solution above 33 0 in rhombohedra. If a glass flask is 
filled with a warm saturated solution of sodium sulphate and the 
neck closed with a loosely-fitting plug of cotton-wool, the solution 
remains clear and liquid after it has been allowed to cool quietly. 
And no change is produced if the plug is carefully removed and 
the end of a glass rod which has been heated in a gas flame 
and allowed to cool is dipped into the liquid. But the slightest 
touch with the other end of the rod, by which particles of dust are 
brought into contact with the liquid, at once causes the crystalliza¬ 
tion of the supersaturated solution, with a corresponding rise in 
temperature due to the liberation of latent heat. 

f nil 

Acid Sodium Sulphate: S 0 2 | 0Na is prepared in the same way 

.as the potassium salt, and is deposited from its warm solution in 
large transparent crystals, without water. At the ordinary tempe¬ 
rature it crystallizes with 1 molecule of water. It decomposes 
when heated into water and sodium disulphate (p. 167), and the 
latter compound is converted into normal sodium sulphate and 
sulphuric anhydride at a higher temperature. 

Sodium Sulphite: So|qn| is obtained from the acid salt by 

adding an equal quantity of sodium carbonate as was used in its 
preparation; it crystallizes without water from its warm solution, 
and with 7 molecules in the cold. 

(OH 

Acid Sodium Sulphite: SO| ONa prepared by saturating a 

solution of sodium carbonate with sulphurous anhydride. Both 
salts are easily soluble in water, and both are largely used by 
brewers for antiseptic purposes—partly to cleanse stale casks, and 
partly as an addition to the beer to prevent acid fermentation. 

Sodium Thiosulphate (Sodium Hyposulphite ): S0 2 j g^ a + 

5 H 2 0 , is obtained by evaporating its aqueous solution in large 
transparent crystals, which melt in their water of crystallization at 


Sodium Nitrate. 369 

56 , and lose all their water at ioo°. At higher temperatures it is 
decomposed. Its preparation and its property of dissolving the 
alogen salts of silver have been already referred to (pp. 168, 169). 
Large quantitie are used for photographic purposes. Owing to 
its decomposition by chlorine, it is also employed to remove the 
last traces of chlorine from articles which have been bleached 
with this substance, and is hence often called an antichlor. A 
cheap mode of preparing it consists in decomposing a solution of 
oxidized alkali-waste (p. 169), which contains calcium thiosulphate, 
with sodium Carbonate; insoluble calcium carbonate and a solu¬ 
tion of sodium thiosulphate are thus obtained. 

Sodium Nitrate (Chili Saltpetre ,): NO.-ONa, crystallizes in 
rhombohedra which so closely resemble cubes that it has been 
called cubic nitre. It is more easily soluble in water than potassium 
nitrate, and only requires about its own weight of water for solution 
at the ordinary temperature. The salt is found in large deposits 
in nature, especially in South America on the borders of Chili and 
Peru, and is an important article of commerce. This crude nitrate 
of soda is, however, far from pure ; it contains sodium chloride, 
iodide and sulphate, together with earthy impurities, and may be 
purified by crystallization. The mother-liquors are used for the 
preparation of iodine (p. 124). Pure sodium nitrate maybe pre¬ 
pared by neutralizing caustic soda or sodium carbonate with nitric 
j acid. 

; It: has already been mentioned that sodium nitrate is slightly 

hygroscopic, so that gunpowder prepared from it becomes moist 
and useless. Nevertheless, sodium nitrate can be readily converted 
into potassium nitrate, and so used indirectly for the preparation 
of gunpowder. It is also used for the manufacture of nitric acid. 

Sodium Carbonate, or Soda : Co|q^ + ioH 2 0 .—Just as 

land-plants require potassium compounds for their growth, so are 
sodium compounds necessary for sea-plants. The former derive 
| their potash from the soil, the latter obtain their necessary soda 
from the sea. And when sea-plants are burnt their ashes contain 
sodium carbonate, in the same way as the ashes of land-plants 
contain the corresponding potassium carbonate. 

Until some eighty or ninety years ago, most of the soda used in 
the arts was obtained in this way from sea-plants, and, at that time, 

13 13 



370 Text-Book of Inorganic Chemistry. 

large quantities of potash were imported into France and Western 
Europe from Russia, Germany, and America. But at the end of 
the preceding century, during the French revolution, this importa¬ 
tion of potash into France almost ceased, and many important 
industries depending upon a supply of alkali seriously diminished. 
A commission was then appointed by the French Government to 
examine into any processes for converting common salt into soda, 
and they decided in favour of the method discovered by Leblanc — 
a process which has remained practically the same from then to 
the present day. Since then this branch of industry has attained 
immense dimensions, and the price of soda has fallen so low that 
in many cases it has displaced the more expensive potash, except 
when this alkali alone can be used— c.g. in the manufacture of 
potash glass, &C. 1 

The quantity of soda which occurs in the mineral kingdom, 
partly in the crystalline state and partly in solution, in the soda- 
lakes of Egypt and North America, is only of local importance, 
and is inappreciably small in comparison with the immense quan¬ 
tities of this salt which are used for so many purposes. 

The conversion of common salt into soda, when casually con¬ 
sidered, does not appear a difficult problem. It might be thought 
that on heating calcium carbonate with sodium chloride the two 
salts would yield sodium carbonate and calcium chloride. But 
even supposing this reaction to take place when the dry sub¬ 
stances are heated, it would not be possible to separate the sodium 
carbonate from the calcium chloride, because, on the addition of 
water, both would be dissolved, and would again produce sodium 
chloride and insoluble calcium carbonate. Such a process could 
only be successful if the calcium chloride were converted into a 
compound insoluble in water. And, as a matter of fact, the process 
introduced by Leblanc depends upon the conversion of a sodium 
salt and calcium carbonate into sodium carbonate and an insoluble 
calcium compound, so that when the mixture is digested with 
water, sodium carbonate alone dissolves. Leblanc used sodium 
sulphate for this purpose, and this salt must first be obtained from 
sodium chloride. 

The process of manufacturing soda, or the alkali manufacture, 
as it is usually called, is not a simple one ; various conditions are 
necessary for carrying it on successfully/and the process consists 
of several distinct chemical operations. The manufacture requires, 
1 See foot-note on p. 349. 


Sodium Carbonate. 37 r 

in the first place, a large supply of common salt and limestone, 
and, secondly, of sulphuric acid. And as the immense quantities 
of sulphuric acid which are necessary for the production of the 
sodium sulphate cannot well be obtained from other manufactories, 
the acid is made on the spot, and sulphuric acid chambers with 
other requisites for manufacturing the acid are one necessity of 
every alkali works. This sulphuric acid is then used in the first or 
salt-cake process , to convert the sodium chloride into sodium 
sulphate or salt-cake ; large quantities of hydrochloric acid being 
then obtained as a bye-product (p. 108). 

After all the preliminary conditions have been fulfilled which 
are necessary to obtain the sodium sulphate or salt-cake on a 
sufficiently large scale, the actual operation of soda manufacture, 
or the conversion of sodium sulphate into sodium carbonate, can 
be commenced. But it is not the object of this small text-book to 
describe all the details of the manufacture of alkali, nor indeed can 
technical chemistry, as little as practical chemistry, be learnt from 
books. A number of practical details, often apparently inex¬ 
plicable, must be attended to if the manufacture is to be a success, 
and these details cannot be learnt from books, but a knowledge of 
them must be acquired in the manufactory itself. And this is true 
not only for the manufacture of soda but also for every other branch 
of chemical industry. There is always a gap between theoretical 
and practical chemistry, which remains unbridged for the man who 
works by rule-of-thumb, but which is readily accounted for by the 
scientific chemist who has learnt to observe accurately and to 
solve chemical problems by experiment. 

It is therefore sufficient if the student understands the chemical 
reactions upon which the change of sodium sulphate into sodium 
carbonate depends. 

In the second process of alkali manufacture—the black-ash 
process —the sodium sulphate (salt-cake) is mixed with calcium 
carbonate (limestone) and small coal, the mixture heated in 
rotating furnaces, and the black mass ( black-ash ) which remains 
behind lixiviated with water to dissolve the sodium carbonate. 
The salt then crystallizes out on evaporating this solution. The 
chemical reactions which produce this change are as follows. 
Sodium sulphate, when heated with the carbon of the coal, is 
converted into sodium sulphide and carbonic oxide :— 

S 0 2 ( 0 Na) 2 + 4C = Na 2 S + 4CO, 


BB2 


372 Text-Book of Inorganic Chemistry. 

and this sodium sulphide is at once decomposed by the calcium 
carbonate, producing sodium carbonate and calcium sulphide :— 
Na 2 S + CO - 0 2 Ca = CO(ONa) 2 + CaS. 

On lixiviating this mass with water, the soluble sodium carbonate 
dissolves, and the insoluble calcium sulphide remains behind. 
This insoluble residue of impure calcium sulphide constitutes 
alkali-waste , and when exposed to the air, gradually forms poly¬ 
sulphides of calcium and other salts which dissolve in the drainage 
water, and which liberate sulphuretted hydrogen when coming 
into contact with acid liquors. Hence, not only is the loss of 
material in the waste very considerable (it contains the whole 
of the sulphur present in the sulphuric acid employed), but it is 
also an intolerable nuisance to the neighbourhood of alkali works. 
Many proposals have been made to recover the sulphur from this 
waste, some of which have proved partially successful. 

In practice, a much larger quantity of limestone is employed 
than that represented by the above equation. This excess of 
limestone becomes converted into quick-lime during the heating, 
and then produces a corresponding quantity of caustic soda, which 
remains in the mother-liquor after the sodium carbonate has 
crystallized out. In some alkali works considerable quantities of 
caustic soda are manufactured in this way. 

Soda is also manufactured to some extent from cryolite, a 
double fluoride of sodium and aluminium : 3NaF, A 1 F 3 (p. 134). 
This compound, when mixed with powdered limestone and strongly 
heated, decomposes into soluble sodium aluminate and insoluble 
calcium fluoride :— 

Na 3 AlF 6 + 3 C 0 - 0 2 Ca = Na 3 A 10 3 + 3CaF 2 + 3CO0. 

And when a stream of carbonic acid is led through the solution 
of sodium aluminate produced on digesting the mass with water, 
sodium carbonate is formed and aluminium hydrate precipitated. 

The direct conversion of common salt into soda has been very 
successfully worked in recent years, and the quantity of soda 
annually produced by this ammonia-soda process is rapidly in¬ 
creasing. It has been found that a mixed solution of common 
salt and acid ammonium carbonate is converted under a pressure 
of about two atmospheres into ammonium chloride and acid 
sodium carbonate :— 

NaCl + Co{qnH 4 = C0|0N a + NH 4 C1. 


Sodium Carbonate. 


373 

The process which is based upon this reaction is carried out 
in the following manner. A saturated solution of common salt is 
mixed with ammonia, and then saturated with carbonic acid under 
a pressure of two atmospheres. The acid sodium carbonate, 
which is only difficultly soluble in water, then separates out, is 
filtered off under pressure, and converted into the normal carbo¬ 
nate by heating. The carbonic acid which is then given off is 
again used to reconvert the ammonia which is recovered from the 
ammonium chloride into ammonium carbonate. 

The ammonia-soda process is not only simpler than the 
older method, but it only yields one bye-product—viz. calcium or 
magnesium chloride, arising from the decomposition of the ammo¬ 
nium chloride by lime or magnesia. Leblanc’s process, on the 
other hand, gives a number of objectionable residues, from which 
only a portion of the sulphur can be again obtained. Still the 
ammonia process can never entirely displace the black-ash process, 
because we cannot dispense with hydrochloric acid which is pro¬ 
duced in such large quantities as a bye-product. 

Although the compounds of sodium and potassium as a general 
rule are very similar to one another, the two carbonates are 
strikingly different, especially in their behaviour with water. 
Potassium carbonate is very soluble in water, and can only be 
obtained in the crystalline form with difficulty; it even attracts 
moisture from the air and deliquesces. Sodium carbonate, on the 
other hand, is much less soluble in water, and can be easily ob¬ 
tained in large crystals, which when exposed to the air rapidly lose 
water and effloresce. A dry fresh crystal placed on the scale-pan 
of a balance loses weight rapidly and continuously. 

The ordinary crystals of sodium carbonate are large trans¬ 
parent monoclinic tablets containing io molecules or 63 per cent, 
of water. When these crystals are heated, they melt at 50° in 
their water of crystallization, to form a clear liquid. At a higher 
temperature water is expelled, and finally the anhydrous salt 
remains behind. Sodium carbonate when crystallized from solu¬ 
tions above 50° contains only 7 molecules of water. The degree 
of solubility of sodium carbonate in water at different temperatures 
varies considerably. The solubility increases with the temperature 
up to about 35 0 , but then begins to diminish, with the separation 
of a salt containing less water. 

The crystals when exposed to the air become covered with a 


374 


Text-Book of Inorganic Chemistry. 


thin, loosely attached layer of a compound containing less water, 
into which the salt is finally completely converted. If a crystal of 
soda which has been long exposed to the air is broken, a small 
nucleus of the transparent undecomposed salt may often be found 
in the centre. 

When strongly heated in a platinum crucible, anhydrous 
sodium carbonate melts to a clear liquid, without undergoing 
decomposition; carbonic acid is not expelled. An aqueous solu¬ 
tion of sodium carbonate, like one of potassium carbonate, reacts 
strongly alkaline. 

Acid Sodium Carbonate : CO —This salt, which is com¬ 

monly known as bicarbonate of soda, is obtained by leading a 
stream of carbonic acid through a strong solution of the normal 
carbonate. The acid carbonate, owing to its smaller solubility, 
then separates out as a crystalline powder. It is usually prepared 
by saturating a mixture of i part of the crystalline normal carbonate 
and 3 parts of the anhydrous salt with carbonic acid. The salt 
crystallizes without water in small distorted monoclinic plates, and 
requires about 11 parts of water for solution at the ordinary tempe¬ 
rature. Its solution reacts neutral, but loses carbonic acid when 
boiled, and is first converted into the so-called sesquicarbonate 
of soda , which is a hydrated double compound of the normal and 
acid carbonates, of the composition :— 


CO 


(ONa 

}ONa 


+ 


2CO 


| ONa 
(OH 


+ 


2H a O. 


This double compound is found in nature as small hard crystals, 
which are permanent in the air. It separates on the banks of the 
soda-lakes of Egypt and America, and is brought into trade under 
the name of Trona. 

Acid sodium carbonate is used to some extent in medicine, 
especially in the preparation of Seidlitz powders. It is found in 
many mineral springs, and particularly in those which are alkaline. 


Phosphates of Sodium.— Of the various compounds of phos¬ 
phoric acid and sodium, the best known and most important is 

Monacid Sodium Phosphate : PO jojq^ 2 + I 2 H 2 0 , which is 

obtainedby nearly neutralizing caustic soda or sodium carbonate with 
ordinary phosphoric acid. On evaporating this solution down, the 


375 


Phosphates of Sodium. 

salt separates out in large transparent crystals which easily effloresce 
in the air. Its solution possesses a faint.alkaline reaction although 
it is an acid salt. The hydrated salt readily loses its 60 per cent, 
of water of crystallization when heated to ioo°; and at a red heat 
it is decomposed into water and sodium pyrophosphate (p. 219). 

Normal Sodium Phosphate'. PO(ONa) 3 +i2H 2 0, crystallizes 
in thin six-sided prisms when a solution of the preceding salt is 
mixed with sufficient caustic soda and evaporated down. The 
solution has a strong alkaline reaction; it absorbs carbonic acid 
from the air, and becomes converted into sodium carbonate and 
the rhonacid phosphate. 

Diacid Sodium Phosphate \ PO |^otT) + ^ 2 0 , is obtained 

when a solution of the monacid phosphate is mixed with as much 
phosphoric acid as it already contains. The solution has an acid 
reaction ; when evaporated down it yields the salt in rhombic 
prisms, which are easily soluble in water. 

Very different from these three salts is the sodium compound 
of monobasic metaphosphoric acid : 

Sodium Metaphosphate : P 0 2 - 0 Na, which is obtained by heat¬ 
ing diacid sodium phosphate :— 

PO {(OH), = PO.-ONa + H 4 0 . 

(p. 219), as an amorphous vitreous mass, and which cannot be ob¬ 
tained in the crystalline form. 

Sodium Pyrophosphate : O j po(ONa) 2 + IO H 2 0 ? is prepared by 

heating the monacid sodium phosphate to redness (p. 219), and sepa¬ 
rates from an aqueous solution of the residue in large crystals, which 
are permanent in the air. Its solution may be boiled without change, 
but if nitric acid is added and the solution then boiled, the salt is 
decomposed into sodium nitrate and diacid sodium phosphate. 

Sodium Arsenate'. AsO+ I2H 2 0, which is isomor- 

phous with the corresponding common sodium phosphate, is de¬ 
posited from its aqueous solution in large well-developed crystals. 

It is largely used in calico-printing and dyeing. The normal 
and diacid arsejiates closely resemble the corresponding phos¬ 
phates. 


376 Text-Bo ok of Inorganic Chemistry . 

Sodium Borate: B 4 0 5 ( 0 Na) 2 + ioH 2 0 (p. 264).—This salt, 
which is commonly known as borax , and is found in nature as the 
mineral tinkal , is prepared by neutralizing a hot solution of sodium 
carbonate with boric acid. When the solution is allowed to 
evaporate slowly the salt separates out as large transparent mono¬ 
clinic prisms of the above composition, which only slightly change 
in the air. At the ordinary temperature 100 parts of water only 
dissolve about 7 parts of the salt, but the same quantity of water 
dissolves more than 200 parts of the salt at ioo°. The solution 
possesses a faint alkaline reaction. From a concentrated aqueous 
solution of borax at 70° or 8o°, the salt crystallizes in octahedra 
with only 5 molecules of water. This octahedral borax is distin¬ 
guished from ordinary prismatic borax by its greater hardness. 
Borax loses its water when heated and swells up into a porous 
mass ; at higher temperatures the salt melts to a dear liquid which 
solidifies on cooling to a hard transparent glass. Fused borax 
possesses the property of dissolving metallic oxides, often in con¬ 
siderable quantities, and the variously coloured glasses which are 
thus produced afford a test for various metals. Thus cobalt oxide 
imparts a blue colour to fused borax ; manganic oxide gives a 
violet, but manganous oxide a colourless glass ; while cupric oxide 
and chromic oxide yield green-coloured glasses. 

Borax is employed in the arts as well as in the laboratory. Its 
property of .dissolving metallic oxides makes it useful in soldering 
two pieces of metal together when the surfaces to be united must 
be perfectly free from oxide. Considerable quantities of borax are 
used for glazing porcelain and in the manufacture of enamels. It 
is also a valuable medicine. 

The compounds of sodium with the halogens and with sulphur 
closely resemble the corresponding compounds of potassium, and 
of these the chloride is the only one which need be referred to in 
detail. 

Sodium Chloride (Common Salt) : NaCl.—This salt not only 
occurs widely distributed in nature, but is also found in immense 
quantities, partly as solid rock-salt , and partly in aqueous solution 
in brine springs and in the sea. 

Common salt crystallizes in the regular system, and generally 
in cubes. It contains no water of crystallization, but usually 
includes small quantities of water mechanically, which cause the 


Sodium CJdoride . 


377 


crystals to decrepitate when heated. It possesses a well-known 
saline taste, and is only slightly more soluble in hot than in cold 
water. 100 parts of water dissolve 36 parts of salt at the ordinary 
temperature, and only 39 parts at ioo°. The concentrated aqueous 
solution, therefore, scarcely contains 27 per cent, of the salt. It is 
insoluble in strong alcohol. When an aqueous solution of common 
salt is evaporated down, the crystallization commences first where 
the concentration is greatest— i.e. on the surface—and the small 
cubes which are thus produced gradually increase in weight, and 
sink to the bottom of the vessel. 

The extraction of salt from the brine springs of Cheshire and 
Worcestershire is an extremely simple process, and consists in eva¬ 
porating the brine in shallow iron pans. In countries where fuel is 
dear and the brine less concentrated— e.g. in Germany, the solution 
is allowed to trickle several times over tall walls of faggots of thorn, 
which exposes a large surface to the air, and causes a rapid evapo¬ 
ration of water. After a concentrated solution has been obtained 
by this process (called graduation ), the strong brine is boiled down 
in iron pans. 

In the same way, salt has also been obtained from sea-water, 
although this liquid only contains about 3-5 per cent, of it. In 
warm southern climates, for example, in the south of France and 
on the coast of Africa, shallow basins are filled by the sea at high 
water, then closed, and the water evaporated by the heat of the sun 
and by the warm winds. Salt obtained from sea-water in its crude 
state, and containing various impurities, is commonly known as 
bay-salt. 

Brine springs are produced by water coming into contact with 
deposits of rock-salt in the earth, and artificial springs are very 
often formed by drilling down to the layer of salt and pumping in 
water. In some cases the rock-salt is mined in the ordinary way, 
and brought to the surface in the solid state. The chief deposits 
of salt occur in Galicia, in various parts of Germany, and at North- 
wich and Droitwich in England. Very important salt deposits have 
been recently discovered at Stassfurt, near Magdeburg, where the 
common salt is associated with potassium chloride and other com¬ 
pounds. 

Rock-salt can be used in its crude state for most technical pur¬ 
poses, but not as a condiment, as it possesses an impure taste due 
to various impurities, which may, however, be separated from it by 
recrystallizing it from its solution in water. 


373 Text-Book of Inorganic Chemistry . 

Commercial common salt often differs both in its external 
appearance and in its taste. In some cases the cubes in which 
it crystallizes are arranged together in large pyramids, in other 
cases they are separate and form a granular crystalline powder. 
One sample may become moist in the air, and possess an intense 
saline taste, while another remains dry, and tastes much less salt. 
The form in which the salt crystallizes depends partly on the 
temperature of evaporation, and partly upon the presence or 
absence of foreign substances, such as sodium sulphate, magnesium 
chloride, or calcium chloride. The two latter salts are those which 
cause the salt to deliquesce in the air, and which impart the sharp 
saline taste that some samples possess. Pure sodium chloride 
does not become moist when exposed to the air, and has a pure 
saline taste. 

Common salt is almost as essential for the life and health of 
men and animals as the air which they breathe, and if a State 
imposes a tax upon this most necessary food (as in Germany), 
while luxuries such as tobacco remain almost untaxed, it proves 
that its rulers lack scientific training, and that they do not under¬ 
stand the simplest and most important questions bearing upon the 
material well-being of the people. 

Detection of Sodium Compounds. 

The intensely yellow colour which sodium compounds impart 
to a Bunsen flame, together with the single yellow line of the 
sodium spectrum (see table of spectra) are so characteristic that 
it is scarcely possible to overlook the presence of even small 
quantities of the metal. Sodium salts may be distinguished from 
those of potassium by the solubility of sodium chlorplatinate : 
Na 2 PtCl 6 , both in water and in alcohol. The two carbonates are 
also very different. Potassium carbonate deliquesces in the air, 
but the sodium salt effloresces. 


LITHIUM. 

Chemical Symbol : Li .—Atomic Weight : 7. 

Lithium is a rare element, but on account of the extensive use 
o some of its compounds in medicine is of considerable value. It 



Lithium . 


379 


is also widely distributed in nature, and is found, though always in 
small quantities, in the minerals : lepidolite, petallite , triphylline , 
&c., as well as in numerous mineral springs, especially in one near 
Redruth in Cornwall. 

The solubility of lithium carbonate in water affords a basis for 
the separation of this salt from the carbonates of the alkaline earths, 
while, on the other hand, its nearly insoluble phosphate permits its 
separation from the easily soluble phosphates of the alkalies. The 
preparation and purification of lithium compounds depend on the 
properties of these two salts. 

Metallic lithium, which may be prepared by electrolysis of its 
fused chloride, is a silver-white tough metal, harder than potassium 
or sodium, but softer than lead. It is distinguished by its low 
specific gravity (0*59), and is, in fact, the lightest solid known. Its 
melting-point is 180° but, unlike potassium and sodium, it is not 
volatile at a red-heat, and can only be distilled at very high tem¬ 
peratures. In the air it oxidizes less easily than potassium and 
sodium, but its freshly-cut surface soon becomes covered with a 
coating of oxide. When heated in the air above its melting-point, 
it catches fire and burns brightly. Water is easily decomposed by 
it at the ordinary temperature, but the hydrogen is not inflamed. 
It burns readily in chlorine, producing its chloride. 

Of the compounds of lithium : lithium oxide and lithium per¬ 
oxide have been but little investigated. 

Lithium Hydrate (. Lithia ): LiOH, which may be prepared by 
boiling the carbonate with slaked-lime, remains behind on evapor¬ 
ating the alkaline solution in a silver dish as a white, transparent, 
easily fusible mass. It resembles caustic soda in external appear¬ 
ance, but is less soluble in water and does not deliquesce in 
the air. 

Lithium Sulphate : S 0 2 joLj + H * 0 ’ cr Y stallizes in rhombic 

prisms and is easily soluble in water .—Lithium Nitrate'. N 0 2 - 0 Li, 
crystallizing without water in rhombohedra, is very easily soluble 
in water, and deliquesces in moist air. 

Lithium Phosphate'. 2PO(OLi) 3 + H 2 0 , is very difficultly solu¬ 
ble in water, and is precipitated on mixing a solution of mon¬ 
acid sodium phosphate with one of a soluble lithium salt as a 
white, heavy crystalline powder, the quantity of which is increased 




380 Text-Book of Inorganic Chemistry . 

on neutralizing the liquid with a little caustic soda or ammonia. 
The salt requires about 2,500 parts of water for its solution at the 
ordinary temperature, and is still less soluble in water containing 
ammonia. 

Lithium Carbogiate : CO | ql!’— com P oun d, which is also 

difficultly soluble in water in comparison with the carbonates of 
potassium and sodium, is deposited as a white powder on adding 
sodium carbonate to a concentrated solution of lithium chloride. 
It dissolves in about 100 parts of water at the ordinary temperature, 
yielding a solution which is faintly alkaline. Hot water dissolves 
more, but it is insoluble in alcohol. Lithium carbonate possesses 
the remarkable property of forming a soluble compound with the 
insoluble uric acid, and is therefore largely taken as a medicine to 
remove this substance from the body, when it separates out during 
the progress of certain diseases (gouty affections, stone, &c.). The 
lithium carbonate is usually given in water containing free carbonic 
acid (lithia-water), in which the salt dissolves more readily than in 
pure water. 

Lithium Chloride : LiCl, crystallizes from very concentrated 
solutions in cubes or octahedra. It is very soluble in water and 
deliquesces in the air. It is also soluble in alcohol, and even in a 
mixture of alcohol and ether. When fused in an open vessel it 
loses some chlorine and becomes partially converted into a basic 
chloride. 

Detection of Xiithium Compounds. 

The smallest quantity of a lithium compound when introduced 
into a non-luminous gas flame gives a splendid crimson colour, the 
spectrum of which is one single line in the red (see table). The 
precipitates produced by sodium carbonate or phosphate in strong 
solutions of lithium compounds also serve to distinguish them from 
those of potassium and sodium. 


RUBIDIUM AND CAESIUM. 

Chemical Symbols : —Atomic Weights : | r 33 *0 

These two rare elements are as yet only of chemical interest. 
Both were discovered by Bunsen and Kirchhoff during their 



Rubidium mid Ccesium. 


381 

researches in spectrum analysis. The two metals are found in 
minute quantities in certain minerals, in the ashes of some plants, 
and in some mineral springs. One of the richest sources of the 
two chlorides is the mineral water of Diirkheim in the Bavarian 
Palatinate, in which caesium was first discovered by Bunsen, but 
even this only contains one part of caesium chloride in 6,000,000 
parts of the water. 

Rubidium , which is obtained, like potassium and sodium, by 
glowing an intimate mixture of its carbonate and charcoal, is a 
white metal with a faint yellowish tint. It is as soft as wax even at 
io°, melts at 38°, and has a specific gravity of 1*5. It at once 
oxidizes when exposed to the air, and then becomes so strongly 
heated that it soon catches fire. When brought into contact with 
water, decomposition at once takes place and the liberated hydro¬ 
gen is ignited. 

The compounds of rubidium so closely resemble those of potas¬ 
sium that it is very difficult to separate the two from one another. 
The double chlorides of the metals with platinum are best adapted 
for this separation, the rubidium compound (2RbCl, PtCl 4 ) being 
much less soluble in water than the corresponding potassium com¬ 
pound. Rubidium is more electro-positive than potassium. Com¬ 
pounds of rubidium when held on a platinum wire in a non- 
luminous gas flame impart a violet colour to it resembling that of 
potassium compounds. The spectrum of this flame is, however, 
quite different, and is characterized by two lines in the red and 
two in the violet. 

Ccesium , which is even rarer than rubidium, has been recently 
obtained in the metallic state by the electrolysis of a mixture of 
caesium and barium cyanides. It is a white, very soft metal, 
melting at 27 0 or a summer temperature, and of 1 *88 specific 
gravity. It decomposes water with explosive violence, and is even 
more electro-positive than rubidium. Caesium is, therefore, the 
most electro-positive of all metals, and has a stronger affinity for 
oxygen, chlorine, &c., than any other metal. Its compounds closely 
resemble those of rubidium and potassium. In order to separate 
it from these two metals, the solubility of its carbonate in absolute 
alcohol, in which the carbonates of rubidium and potassium do not 
dissolve, may be made use of. The spectrum of caesium compounds 
is especially characterized by two lines in the bright blue. 


382 


Text-Book of Inorganic Chemistry . 


AMMONIUM. 

It has already been remarked (p. 188) that we are not acquainted 
with the substance called ammonium in the free state ; but the fact 
that it forms an amalgam with mercury and the close relation of its 
compounds to those of potassium, justify us in deciding in favour 
of its metallic nature, and in placing it together with the metals of 
the alkalies. Its composition is expressed by the formula : NH 4 , 
or perhaps more correctly as N 2 H 8 = (NH 4 ) 2 . 

Ammonium oxide and hydrate are just as little known as 
ammonium itself. It has sometimes been stated that the aqueous 
solution of ammonia in water contains ammonium hydrate : 
NH 3 + H ,,0 = NH 4 0 H, but this is as little true as the assertion 
that a solution of carbonic anhydride in water contains carbonic 
(OH 

acid : CO - The mere fact that it is only necessary to blow 

a current of air or some other indifferent gas through these solutions 
to expel the whole of the gas, sufficiently proves that they are not 
true chemical compounds. 

The ammonium salts are easily obtained by neutralizing 
ammonia or ammonium carbonate with the corresponding acid. 
Like those of potassium, they are nearly all soluble in water, and 
mostly crystallize well. 


Ammonium Sulphate : S0 2 1 This salt, which is iso- 

morphous with potassium sulphate, can easily be obtained in the 
crystalline state. It melts at 140°, and decomposes at higher tem¬ 
peratures into ammonia, water, nitrogen, and ammonium sulphite, 
which sublimes with some unchanged ammonium sulphate. It is 
manufactured on a large scale by leading the ammonia evolved on 
heating the ammoniacal-liquor of the gas works with milk of lime 
into dilute sulphuric acid until saturation, and then evaporating 

down. An acid salt, of the composition : S0 2 j is also known. 

Besides these compounds, normal salts, in which one of the atoms 
of ammonium in ammonium sulphate is displaced by potassium or 


Ammonium Carbonate. 


333 


sodium, also exist; they have the composition : S 0 2 


SO 


ONa 

nho 4 


+ 2H20. 1 


OK 

ONH 4 


and 


Ammonium Nitrate: NO.,-ONH 4 , crystallizes in six-sided 
prisms, without water, and is isomorphous with potassium nitrate. 
It is easily soluble in water, absorbing heat, and deliquesces in the 
air. When heated it first melts and is then converted into nitrous 
oxide and water (p. 198). 


Ammonium Phosphates.— The compounds of phosphoric acid 
with ammonia closely resemble those of sodium and potassium. 

A salt corresponding to the monacid phosphate: POj 

in which one atom of ammonium is displaced by one of sodium, is 
the compound called microcosmic salt , or monacid sodium-ammo- 
(OH 


nium phosphate : PO -j ONH 4 + 4FCO. It is easily obtained by dis- 
(ONa 


solving 6 or 7 parts of common sodium phosphate and 1 part of 
ammonium chloride in 2 parts of boiling water, and separates out 
on cooling in well-defined crystals. It may be freed from the 
sodium chloride which it contains by recrystallizing with the addi¬ 
tion of a little ammonia. 

This salt loses its water when heated, and then gives off am¬ 
monia, leaving sodium metaphosphate: P 0 2 - 0 Na, as an easily 
fusible mass. For this reason the salt is used in blowpipe analysis. 
A fragment heated on a platinum wire in a Bunsen burner gives a 
glass bead of sodium metaphosphate, which, like borax, can dissolve 
various metallic oxides and become coloured by them. Silica is 
not dissolved by the fused salt, and may thus be readily detected. 


Ammonium Carbonate.— On the union of ammonia and 
carbonic acid, different compounds are produced according to the 
conditions under which the combination takes place. It has 
already been stated (p. 290) that if the two gases act upon one 

another in the absence of water, ammonium carbamate : CO j qnH 4 
1 These may also be regarded as double salts of the two sulphates, thus 
so, {g£+SO, {g£% and SO, {g^+SO, {g^+^O, but are best 
represented by the above formulae.— Ed. 


384 Text-Book of Inorganic Chemistry . 

and not ammonium carbonate is formed. The same salt is also 
produced when aqueous ammonia is saturated with carbonic acid, 
although the chief product of this reaction is ammonium carbo¬ 
nate. And it is highly probable that the normal as well as the 
acid salt which crystallize from concentrated solutions (the former 
when ammonia is in excess, the latter with excess of carbonic acid) 
contain ammonium carbamate. 

The commercial ammonium carbonate (carbonate of ammonia) 
is manufactured by heating a mixture of ammonium chloride or 
sulphate with powdered chalk ; it then sublimes as a translucent, 
crystalline solid, with a strong ammoniacal odour. When exposed 
to the air, it crumbles to a white crystalline powder, consisting 
chiefly of acid ammonium carbonate. This substance was pre¬ 
viously thought to be a mixture of the normal and acid ammonium 

caibonates, of the composition : CO | ON + ^ {oNH ^ ence 

its old name of sesquicarbonate, but is now known to contain 
considerable quantities of ammonium carbamate, and is therefore 
a mixture of several salts. The presence of the ammonium 
carbamate may be readily shown by shaking the freshly prepared 
solution with calcium chloride, and filtering off the clear solution 
from the precipitated calcium carbonate. This solution contains 
calcium carbamate, which gradually decomposes on standing, 
more quickly if heated, with the formation of insoluble calcium 
carbonate. A solution of ammonium carbonate free from the 
carbamate does not give the same result; nor does a solution of 
the commercial salt which has been prepared for some time, since 
ammonium carbamate when dissolved in water gradually, or quickly 
if warmed, takes up the elements of water, and becomes converted 
into ammonium carbonate. 

Ammonium carbonate is produced during the destructive dis¬ 
tillation of nitrogenous organic substance—^, coal. And although 
the quantity of nitrogen contained in coal is very small, and only 
a part of this nitrogen is converted into ammonia, still the immense 
quantities of coal which are used for the manufacture of coal-gas 
are so great that very nearly all the ammonium compounds now 
brought into trade are derived from the ammoniacal liquors of the 
gas-works. 

Ammonium carbonate is also produced by the putrefaction of 
organic compounds containing nitrogen. Large quantities are 
always present in putrid urine and sewage, and are produced from 


Ammonium Chloride . 


38s 


the decomposition of the urea which they contain. This compound 
takes up two molecules of water during putrefaction, and becomes 
converted into ammonium carbonate :— 



CH 4 N 2 0 + 2 H 2 0 


Urea 


Ammonium Chloride ( Sal-ammoniac ) : NH 4 C 1 .—This salt is 
brought into trade in two forms. First, when crystallized from 
water as a white powder of small octahedra, or cubes, grouped 
together; and secondly, when obtained by sublimation as a fibrous, 
crystalline, and compact mass. It possesses a saline taste, and is 
soluble in about 2-5 parts of water at the ordinary temperature, or 
in rather more than its own weight at ioo°. It is nearly insoluble 
in alcohol. When heated it volatilizes, without melting and with¬ 
out decomposition. At a higher temperature its vapour dissociates 
into hydrochloric acid and ammonia, which reunite on cooling. 

In former times large quantities of sal-ammoniac were imported 
from Egypt, where, for lack of other fuel, dried camel’s dung was 
burnt. The nitrogen and common salt contained in the dung gave 
rise on burning to ammonium chloride, which was deposited with 
the soot in the chimneys, and was afterwards purified by some 
crude method. 

Beautiful crystals of ammonium chloride are sometimes found 
in the cavities of solidified lava, and it was formerly thought that 
the salt had been sublimed with the lava from the interior of the 
earth. But it has been proved by Bunsen, from observations of 
volcanic phenomena in Iceland, that crystals of ammonium chloride 
are only found where the glowing lava has flowed over ground 
covered with vegetation. The ammonia, produced by the dry dis¬ 
tillation of the plants would then unite with the free hydrochloric 
acid contained in the lava, and so produce ammonium chloride. 

Ammonium chloride, like all other ammonium compounds, is 
now almost exclusively obtained from the ammoniacal liquor of the 
gas-works (p. 322). The ammonia which is evolved on heating 
the gas-liquor with milk of lime is passed into dilute hydrochloric 
acid until saturated, and the solution so obtained then evaporated 
down. The crude salt is afterwards purified, either by crystalliza¬ 
tion or by sublimation. 

Sal-ammoniac is largely used in the arts for the preparation of 
ammonia (p. 184), as a valuable medicine, in dyeing, and for many 
other purposes. 


c c 


386 Text-Book of Inorganic Chemistry. 

Ammonium Bromide'. NH 4 Br, closely resembles the chloride, 
but is still more soluble in water. It is largely used in photography 
and in medicine. 


Ammonium Sulphide : (NH 4 ) 2 S, and Ammonium Sulphy- 
drate : (NH 4 )HS.—Both compounds may be obtained in the 

crystalline form by allowing dry ammonia and dry sulphuretted 
hydrogen to act upon one another at a low temperature (— 15 0 ). 
When the ammonia is in excess the former compound is produced, 
with excess of the sulphuretted hydrogen, the latter. Ammonium 
sulphydrate may also be prepared by leading sulphuretted hydrogen 
into a saturated solution of ammonia in absolute alcohol, when the 
compound crystallizes out. The two compounds are usually pre¬ 
pared in aqueous solution. If aqueous ammonia is completely 
saturated with sulphuretted hydrogen, ammonium sulphydrate is 
formed, and if to this liquid the same quantity of ammonia is added 
as was used for its preparation, it is converted into ammonium 
sulphide. 

Both are colourless liquids, soon becoming yellow when exposed 
to the air, and smelling both of ammonia and sulphuretted hydro¬ 
gen. The yellow colour which the solutions acquire when exposed 
to the air is due to partial oxidation of the sulphuretted hydrogen 
which they contain. During this process water and free sulphur 
are produced, and the latter substance then dissolves in the unde¬ 
composed sulphide. Ammonium sulphide, like potassium sulphide, 
can dissolve large quantities of sulphur, and so produce polysul¬ 
phides corresponding to those of potassium. 

Ammonium sulphide is much used in analytical chemistry, 
partly to precipitate the sulphides of those metals which are 
insoluble in water, but which are not formed by the action of 
sulphuretted hydrogen on an acid solution, and partly to dis¬ 
solve those insoluble sulphides which have the character of 
sulpho-acids (eg. antimony sulphide). With such sulphides it 
forms soluble sulphosalts, and they may be thus separated from 
other insoluble sulphides (eg. lead sulphide) which have not this 
property. 


Oxyammonium Salts.— Just as ammonia unites with acids and 
forms the ammonium salts, in the same manner the derivative of 
ammonia—oxyammonia or hydroxylamine : NH 2 -OH (p. 186)— 


Detection of Ammonium Compounds. 387 

produces with acids the salts of the radical oxyammonium. The 
radical itself: NH 3 OH or (NH 3 OH) 2 , is as little known in the free 
state as is ammonium. 

Oxyammonium Sulphate-. SO,{ggJH,OH) crystallizes from 

its aqueous solution in large colourless prisms, and is insoluble in 
alcohol. 

Oxyammonium Chloride : NH 3 OH-Cl, is soluble both in water 
and in absolute alcohol, and by this latter property can be separated 
from ammonium chloride. It crystallizes from water in tablets, 
from alcohol in monoclinic prisms, melts at 151 0 , and is decom¬ 
posed at a slightly higher temperature into nitrogen, hydrochloric 
acid, ammonium chloride, and water. 


Detection of Ammonium Compounds, 

The similarity between the ammonium and potassium com¬ 
pounds is so close that on mixing solutions of ammonium chloride 
and platinum chloride a yellow crystalline precipitate of ammonium 
chlorplatinate : (NH 4 ) 2 PtCl 6 , is produced, just as difficultly soluble 
in water, and just as insoluble in alcohol as the potassium com¬ 
pound. If, however, the ammonium compound is heated, ammo¬ 
nium chloride and chlorine are evolved, and pure spongy platinum 
remains behind, from which water extracts nothing. 

Compounds of ammonium are best recognized by their beha¬ 
viour with caustic alkalies (potash, soda, or lime), which liberate 
ammonia even in the cold, but more readily when gently warmed. 
This free ammonia can be detected by its odour, and by the white 
fumes of ammonium chloride produced when a glass rod, moistened 
with hydrochloric acid, is brought near the liquid. 

Minute traces of free ammonia or ammonium compounds may be 
detected by Nesslei s solution an alkaline solution of potassium- 
mercuric iodide. When a few drops of this solution are added to 
a dilute solution containing ammonia or a salt of ammonium, a 
reddish-brown precipitate or corresponding colour is at once pro¬ 
duced. . In this way it is possible to detect 3 ^ of a milligramme of 
ammonium chloride dissolved in 50 c.c. of water. 


c c 2 



3 88 


Text-Book of Inorganic Chemistry. 


METALS OF THE ALKALINE EARTHS. 

The four metals belonging to this group : calcium ,, strontium, 
barium, and magnesium, are distinguished from the metals of the 
alkalies by their weaker attraction for oxygen, and further by the 
fact that their oxides, or rather hydrates, are much less soluble 
than those of the alkali-metals, and do not deliquesce in the air. 
Their sulphates (except that of magnesium) are also much less 
soluble in water than the alkaline sulphates, and their carbonates 

are quite insoluble. . . 

By far the most important of this group of metals is calcium, 

the one we have placed first. 


CALCIUM. 

Chemical Symbol'. Ca.— Atomic Weight'. 40. 

Calcium, with oxygen, silicon, and aluminium, belongs to those 
elements which are most widely distributed in nature. It is never 
found in the metallic state, but usually in combination with oxygen 
as some salt, never as the free oxide. In combination with fluorine 
it forms the important mineral fluor-spar, and occurs as its chloride 
in many mineral waters and in the sea. 

The most important oxygen compounds of calcium found in 
nature are the very various forms of its carbonate, its sulphate (as 
gypsum), its phosphate, and its silicate, which is a constituent of 
many rocks. 

Metallic calcium, which is difficult to prepare pure, is only of 
theoretical interest. It may be obtained by heating calcium iodide 
with sodium in a well-closed iron crucible, or by heating a mixture 
of 3 parts of calcium chloride, 4 parts of zinc, and 1 part of sodium 
in a Hessian crucible. An alloy of zinc and calcium is thus 
obtained, from which the zinc may be volatilized by strong heating 
in a crucible of gas-carbon, placed inside a larger one of clay, then 
leaving a regulus of metallic calcium behind. The best method of 
all of obtaining pure calcium is by the electrolysis of its fused 

chloride. , r , , , , 

Calcium is a brilliant pale yellow metal, of about the same hard- 



Calcium. 


389 

ness as lead. Its specific gravity is i*6 ; it is malleable, but 
becomes brittle when hammered. When exposed to moist air, 
calcium soon becomes covered with a grey coating of oxide and is 
gradually converted into calcium hydrate. It decomposes water 
at the ordinary temperature, but the hydrogen does not catch fire. 
The metal burns brilliantly when heated to redness in the air, or in 
chlorine. Dilute nitric acid acts violently on it, but the concen¬ 
trated acid does not attack it. 

Calcium, like the other metals of the group, is nearly always a 
dyad element; it appears to possess a valency of four in the per¬ 
oxide, and possibly also in bleaching-powder. 

Calcium Oxide, Quick-lime : CaO.—This oxide can be far 
more easily obtained than the oxide of potassium or sodium. 
It suffices to heat the carbonate strongly in the air in order to expel 
the whole of the carbonic acid, and leave the oxide behind. The 
product so obtained is then more or less pure according to the 
purity of the material used for its preparation. Crystallized calcite, 
especially transparent Iceland-spar, yields oxide chemically pure, 
and white marble nearly so. 

It is remarkable that calcium carbonate is only completely de¬ 
composed by heating when a second indifferent gas is present. If 
heated in an atmosphere of carbonic acid, decomposition goes on 
until a certain pressure (depending on the temperature) is reached, 
and then ceases. Marble or Iceland-spar when heated in a closed 
crucible is only partially decomposed, because it is always sur¬ 
rounded with an atmosphere of carbonic acid. Complete decom¬ 
position can, however, be easily brought about by placing a piece 
of charcoal on the bottom of the crucible. The glowing charcoal 
then converts the carbonic acid into carbonic oxide, into which the 
former gas can diffuse until all the calcium carbonate is decom¬ 
posed. 

Calcium oxide or quick-lime is a white amorphous mass, usually 
retaining the form of the calcium carbonate from which it has been 
prepared. It is infusible, even in the oxy-hydrogen blowpipe, but 
becomes then so strongly heated as to emit a most intense white 
light (lime-light). When exposed to the air, it absorbs water and 
carbonic acid, and crumbles to a fine white powder, which effer¬ 
vesces with hydrochloric acid. 

If water is poured on to quick-lime, the porous mass first absorbs 
considerable quantities of the liquid like a sponge. After a short 


390 


Text-Book of Inorganic Chemistry. 

time chemical action sets in, and the calcium oxide unites with 
water to form calcium hydrate :— 

CaO + H „0 = Ca(OH) 2 . 

The reaction is accompanied with so large an evolution of heat that 
a portion of the water is converted into steam. At the same time 
the piece of quick-lime swells up and is finally converted into a dry, 
white, soft powder of the following compound, calcium hydrate. 


I OH 

Calcium Hydrate, Slaked-lime : Ca- Qpj—This hydrate is 


slightly soluble in water, i part requiring 600 parts of cold water 
to dissolve it. In hot water it is less soluble than in cold, whence 
it follows that a cold saturated solution when boiled becomes turbid, 
and calcium hydrate separates in the crystalline form, but redis¬ 
solves as the solution again cools. The solution has a distinct 
alkaline reaction. 

Slaked-lime mixed with water forms a turbid solution of the 
hydrate with undissolved suspended particles. Such a mixture is 
called milk of lime. If this is filtered, the suspended particles 
remain behind on the filter, and clear alkaline lime-water runs 
through. Lime-water must be preserved in well-stoppered bottles, 
as it absorbs carbonic acid from the air and becomes turbid from 
the insoluble calcium carbonate so produced. This property of 
lime-water makes it useful for detecting carbonic acid, even when 
mixed with large quantities of other gases— e.g. in coal-gas. 

Calcium hydrate easily loses its water when heated to redness 
—more easily than calcium carbonate parts with its carbonic acid— 
and is reconverted into calcium oxide. 


Slaked-lime is largely employed for numerous and highly im¬ 
portant technical purposes. Immense quantities are used for the 
preparation of mortar. For this purpose it has been employed for 
thousands of years, and it may probably be considered as one 
of the first chemical compounds artificially prepared by man. 
It would be interesting to know how and when the knowledge 
was acquired that limestone, when heated to redness, becomes 
changed into a new substance, and that this quick-lime when 
made into a paste with water gradually hardens in the air, or if 
placed between stones, cements them together. But questions of 
this nature must, for the present, remain unanswered, since the 
preparation of mortar from lime was known even in pre-historic 
times. 


Calcium Hydrate . 391 

Ordinary mortar is prepared by mixing slaked-lime and sand 
together with water to form a paste, which is spread between the 
bricks or stones to be cemented together. The hardening of the 
mortar only takes place slowly in the course of years—many years 
if the walls are thick, and is the more complete the older the walls. 
And not only does the mortar itself gradually become hard, but it 
also attaches itself firmly to the stones between which it is placed, 
and it often happens when old buildings are pulled down that the 
stones themselves break more easily than the mortar which holds 
them together. The same effect may be noticed in some of the 
calcareous conglomerates or pudding-stones, which often consist 
of quartz pebbles cemented together with limestone, and which if 
broken by the blow of a hammer usually fracture across the pebbles 
and not at the joints. 

The hardening of mortar is due to the action of the carbonic 
acid of the air, which gradually unites with the calcium hydrate to 
form calcium carbonate, while water is expelled :— 

CaJ 2 S + CO„ = COO.Ca + H..O. 

(Uri 

The sand which the mortar also contains serves to make the 
mass more porous, so that the carbonic acid can penetrate into the 
interior of the walls; it also causes the mortar to adhere more 
firmly to the bricks or stones. That old buildings—^, those of 
the Romans or of the Middle Ages—possess greater solidity than 
our modern structures is not only because a better kind of mortar 
was used, but is a natural consequence of the longer action of the 
carbonic acid. A mortar which has only been exposed to the air 
for a few years, and which still contains unchanged calcium 
hydrate, is of course less hard than one on which the carbonic acid 
has acted for centuries, and in which the whole of the calcium 
hydrate has been converted into calcium carbonate. 

It is considered unhealthy to live, and especially to sleep in 
the rooms of newly built houses, because the moisture which is 
contained in the walls continually saturates the air of the rooms. 
The water mechanically mixed with the mortar soon evaporates 
from the porous walls, but the walls still remain moist from the 
water which is gradually set free chemically during the union of 
the calcium hydrate with carbonic acid. The hardening of the 
mortar, and consequent drying of the walls, can be accelerated by 
burning coke or charcoal in open grates in the closed rooms. The 


392 Text-Book of Inorganic Chemistry. 

large quantities of carbonic acid so produced then quickly pene¬ 
trate into the walls, and the water which is set free is evaporated 
by the heat; and to allow this water vapour to escape, it is 
necessary to open the doors and windows occasionally. In this 
way the walls of new buildings may be dried in as many hours as 
they would otherwise require weeks or even months. 

The manufacture of the large quantities of quick-lime which 
are used for the preparation of mortar is carried on in furnaces 
specially constructed for the purpose—called lime-kilns. Two 
forms of these kilns are shown in figs. 61 and 62. Both are built 
of firebricks in the form shown. In the one kiln (fig. 61), which 
is but little used, large pieces of limestone are built below in the 



Fig. 61. Fig. 62. 


form of an arch, and the whole kiln filled up with the same material, 
leaving free spaces for the hot gases to pass through the entire 
mass. A fire is then made in the arch below, and continued until 
the whole of the limestone has been converted into quick-lime. 
After cooling, the lime is extracted, and the process repeated. 
Fig. 62 represents a much more economical form, as it is con¬ 
tinuous in its action* The kiln is filled with pieces of limestone, 
and a fire kept up continuously at the side, a. The lime is raked 
out from the bottom of the kiln as it is produced, and fresh lime¬ 
stone introduced at the top to supply its place. A form of kiln 
largely used in England is similar in shape to fig. 62, but without 
a side opening for the fire. The kiln is filled with alternate layers 
of limestone and small coal, and the lowest layer of coal ignited. 











393 


Calcium Hydrate. 

As the coal gradually burns away, the lime is abstracted from the 
base, add fresh alternate layers of limestone and coal are intro¬ 
duced from the top. 

Some varieties of limestone contain other compounds than 
calcium carbonate, such as magnesium carbonate, alumina, and 
silica. If these are burnt, a lime is obtained which only becomes 
slightly heated when slaked, and which is unsuitable for the pre¬ 
paration of ordinary mortar. In the case of an argillaceous 
limestone— i.e. one mixed with clay (aluminium silicate), the lime 
decomposes some of the clay when moistened with water, producing 
compounds of lime and alumina and of lime and silica, which are 
insoluble in water, and which, in contact with water, gradually 
become hard and solid. If such an argillaceous limestone or a 
mixture of limestone and clay is burnt, then powdered and mixed 
with water, it gradually hardens, owing to the formation of these 
insoluble compounds, and becomes harder and harder the longer 
it remains in contact with water. Mortars of this description, 
known as hydraulic mortars, are therefore especially suited for 
buildings which are always covered with water. 

Hydraulic mortars are prepared from various materials and in 
varying proportions—some hardening slowly and some quickly. 
The rapidity with which the mortar hardens depends upon the propor¬ 
tion of clay present : thus Roman cement , containing about 30 per 
cent, of clay, hardens in a few hours ; while Portland cement, largely 
made in England by mixing chalk with clay and then carefully 
burning, contains less clay and hardens more slowly. A hydraulic 
mortar may also be prepared by mixing quick-lime with soluble 
silicic acid or with an aluminium silicate which has been acted on 
by hydrochloric acid ; such a silicate is decomposed by the lime in 
the presence of water. 

Calcium hydrate is also used for a large number of other tech¬ 
nical purposes—for example, for the preparation of ammonia from 
ammonium chloride, for the manufacture of the caustic alkalies and 
of bleaching powder, for the saponification of fats, for the purifica¬ 
tion of coal-gas, and the manufacture of hard glass, in tanning to 
remove hair from the skins, &c. 

Besides calcium oxide, a second, far less important oxygen 
compound of calcium is known, viz. 

Calcium peroxide : CaO, 2 , which cannot, like barium peroxide, 
be prepared by heating calcium oxide in oxygen. It is, however, 


394 Text-Book of Inorganic Chemistry. 

easily obtained by adding hydrogen peroxide to lime-water, when 
it is precipitated in crystalline plates. When gently heated it 
readily loses one-half of its oxygen. 

Of the other compounds of calcium, the following are of most 
general and chemical importance. 

Calcium Sulphate: S 0 2 - 0 2 Ca.—This slightly soluble salt is 
largely distributed in nature in different forms, and usually accom¬ 
panies common salt. The anhydrous compound, of the above 
composition, is known as anhydrite , and occurs either in crystals 
belonging to the rhombic system or more generally in granular or 
fibrous crystalline masses. 

Far commoner than this is the compound of calcium sulphate 
with two molecules of water : S 0 2 • 0 2 Ca + 2 H 2 0 , which in its most 
general form is called gypsum. This compound is found in nature 
in very various modifications, but usually as a compact crystalline 
mass, generally of a greyish colour : gypsum proper or common 
gypsum. Colourless, transparent crystals are also often found : they 
crystallize in the monoclinic system, and twins are common. A 
granular, crystalline pure variety of gypsum, highly valued as an 
ornamental stone, is known as alabaster. Finally gypsum is often 
found as fibrous masses : fibrous gypsum ; or in transparent plates, 
which very easily split into thin leaves, as selenice. Selenite is 
largely used for optical purposes. Gypsum may be readily pre¬ 
pared artificially by mixing a solution of calcium chloride with 
sulphuric acid or a sulphate, when it falls as a crystalline preci¬ 
pitate. 

Crystalline gypsum is so soft that it can be scratched with the 
nail. As mentioned above, it is only slightly soluble in water, either 
hot or cold—one part requires nearly 400 parts of water for solution. 
The solubility is increased by the addition of common salt, so that 
one part of gypsum dissolves in about 120 parts of a saturated 
solution of common salt. 

The two molecules of water which gypsum contains are only 
loosely united to it. It parts with the greater portion when heated 
up to 120 0 , but the residue is only expelled above 200°. Gypsum 
which has been heated to a temperature not much above 120° 

(Plaster of Paris ) acquires the property of again uniting with water 
when mixed with it, and of setting to a hard mass. Plaster of 
Paris, prepared as above, and ground to a fine powder, is largely 
used for various technical purposes. When mixed with water it 


Calcium Sulphate. 395 

soon sets hard, and can therefore be employed as a cement, which 
will not, however, withstand the action of water. As the hardening 
(with an evolution of heat) sets in soon after the plaster is mixed 
with water, only a small quantity must be prepared at a time; but 
if a strong solution of common alum is used instead of water the 
setting takes place more slowly, and so allows more time for mani¬ 
pulation. If a stronger cement is required, the plaster of Paris is 
first mixed with iron filings and then with water. Such a mixture 
is sometimes used to cement the iron parts of a building into the 
walls. 

It is a well-known, but remarkable fact, that if gypsum is 
heated considerably above I20°it loses the property of uniting 
rapidly with water. And in the same manner, the naturally occur¬ 
ring anhydrous calcium sulphate (anhydrite) cannot unite with 
water to form a hard mass. Both forms combine, however, slowly 
with water when left in contact with it for a long time, and finally 
produce ordinary gypsum. 

As the mixture of plaster of Paris and water readily adapts 
itself to any surface with which it is brought into contact, it is 
used for producing copies of coins and medals, as well as of busts 
and statuettes (plaster casts).. It was also previously used in the 
process of stereotyping, in which a cast is taken of the type, and from 
this a second cast with fusible metal, which can then be used for 
printing just as the original type. Plaster casts for this purpose 
have now been to a large extent superseded by paper pulp, which 
has the special advantage that it is not so brittle as the dry plaster. 
If plaster of Paris is mixed with a solution of glue or gelatine in¬ 
stead of with water, it solidifies more slowly, and is hard enough 
to take a good polish. The mixture may be coloured with various 
metallic oxides, and is used to prepare artificial marbles, &c. 

Large quantities of plaster of Paris are used to coat the interior 
walls and ceilings of our rooms, and it is also employed in medicine 
and agriculture. The surgeon uses it for the preparation of plaster 
bandages, which preserve a joint surrounded with them from move¬ 
ment ; and the agriculturist employes it to impart ammonia from 
the air to plants sprinkled with it. Plaster of Paris has the power 
of absorbing ammonia, and attracts this substance when exposed to 
the air, yielding it again slowly to the plants with which it is in 
contact. Naturally occurring gypsum and anhydrite do not possess 
this property. 


396 Text-Book of Inorganic Chemistry. 

Calcium Nitrate-. ^q-] 0 2 Ca + 4H 2 0, may be easily obtained 

by neutralizing nitric acid with calcium carbonate, and crystallizes 
from its aqueous solution when evaporated down to a syrup. It is 
easily soluble both in water and alcohol. The salt, when freed 
from its water of crystallization by heating, forms a porous white 
deliquescent mass. Calcium nitrate is produced when lime— e.g. 
that contained in the mortar of walls—is exposed to the vapours 
of ammonia, which, in the presence of the strong base, is then 
oxidized to nitric acid by the oxygen of the air. 

In buildings where considerable quantities of ammonia are pro¬ 
duced—^. in stables and cow-stalls—the quantity of calcium nitrate 
accumulating in the walls may often be sufficient to crystallize out 
in dry weather, and its deliquescent nature is one reason why the 
walls of such buildings are usually moist. Even the human body 
emits small quantities of ammonia, which is also absorbed in the 
walls and converted into calcium nitrate. For this reason the 
rubbish of old walls when extracted with water often yields con¬ 
siderable quantities of calcium nitrate. 


PO) 

Calcium Phosphate: Ca 3 P 2 0 8 , or p Q [ 0 6 Ca 3 .—This salt is 

widely distributed in nature, and is sometimes found in consider¬ 
able quantities ; and since its immense importance in agriculture 
has been recognized, new localities of its occurrence have been 
and probably will be discovered. 

To the minerals which consist essentially of calcium phosphate 
belong phosphorite , found in compact masses in some parts of 
Spain, Germany, and other places, and sombrerite , imported in 
large quantities from some of the islands of the Antilles, particu¬ 
larly Sombrero. In combination with calcium fluoride and chloride, 
it forms the mineral apatite : 3(PO) 2 O e Ca 3 + Ca(F,Cl) 2 , which is 
found either crystalline or in a massive form. Finally calcium 
phosphate is the chief mineral constituent of bones, which when 
burnt leave a residue of calcium phosphate and carbonate, bone- 
ash. The excrements of many animals, especially of those which 
are carnivorous —coprolites (petrified droppings), urine, guano, &c. 
—are all more or less rich in calcium phosphate, and are therefore 
valuable manures. 

The calcium phosphate which we require for the growth of our 


Calcium Phosphate. 39 7 

bodies, principally for the formation of bone, we derive directly and 
indirectly from plants, and these again, to whom it is just as neces¬ 
sary, absorb it by their roots from the soil. Calcium phosphate is 
as necessary as carbonic acid for the growth of most plants, par¬ 
ticularly cereals. A grain of wheat, barley, or rye, planted in a 
soil absolutely free from phosphates, will, it is true, germinate and 
grow to a certain stage, but will produce no fruit. And if we sow 
cereals year after year on the same plot of ground, and so gradually 
make it poorer and poorer in calcium phosphate without returning 
any to the soil, the crops become less and less, and finally the 
ground is unfruitful. 

For centuries it has been known to the farmer that he can pre¬ 
serve the fertility of his land by the addition of animal excrements, 
without suspecting what was the action of these substances. But 
Liebig, the founder of scientific agriculture, showed that this 
manure again gives to the land the mineral constituents (amongst 
other things) which have been abstracted from it by the harvest, 
and that the productiveness of the soil can be still further increased 
by the addition of those mineral substances which the plants re¬ 
quire for vigorous growth. 

The scientific agriculturist is, therefore, no longer satisfied 
with farmyard manure (which is indeed often insufficient in 
quantity), but adds to his land artificial manures, prepared in 
chemical factories from suitable materials. To these substances 
belongs calcium phosphate, and this compound, in the form of 
bone-meal, or of minerals containing it, has been successfully 
added to soils originally poor in calcium phosphate or from which 
this substance has been gradually abstracted by continuous har¬ 
vests. Calcium phosphate is, however, insoluble in water, and is 
only slowly dissolved by water containing carbonic acid, hence it 
can be only slowly absorbed by the roots of the plants. A soluble 
calcium phosphate acts more quickly ; and the farmer requires the 
phosphate in this form if he is to materially increase the fertility of 
his land The chemical process by which this change is effected 
is a very simple one. Phosphorite, bone-ash, or some other min¬ 
eral containing calcium phosphate is powdered and treated with 
about a half or two-thirds of its weight of sulphuric acid. A por¬ 
tion of the calcium is then converted into calcium sulphate, setting 
free some phosphoric acid which forms a soluble acid salt with the 
undecomposed calcium phosphate : 


398 


Text-Book of Inorganic Chemistry. 


PO 10 «Ca 3 + SO a (OH) 2 


2 S 0 3 - 0 2 Ca + 


pouoh) 4 

PO J 0 . 2 Ca 


The mixture of calcium sulphate and acid calcium phosphate is 
not separated, but is brought in this form as a manure into trade, 
and is called superphosphate. The acid calcium phosphate is 
soluble in water, and when mixed with the soil is at once absorbed 
by the roots of the plants. 

Chemically pure normal calcium phosphate can be obtained by 
mixing a solution of calcium chloride with one of normal sodium 
phosphate, or with the common acid phosphate if a little ammonia 
is afterwards added. The gelatinous precipitate dries to a white 
earthy powder. It is insoluble in water, but soluble in hydrochloric 
acid, from which solution it is again precipitated unchanged by 
ammonia. 


Monacid Calcium Phosphate : HCaP 0 4 , or PO | Qp^ a is preci¬ 
pitated as a gelatinous mass, often becoming crystalline on stand¬ 
ing, when a solution of calcium chloride is mixed with one of 
monacid sodium phosphate and acetic acid :—• 

PO ioH a)2 + CaCI * “ PO {oH a + 2NaCL 

Like the normal salt, it is insoluble in water. 


Diacid Calcium Phosphate'. H 4 CaP. 2 0 8 , or pq } o'ca 4 may 

obtained from the normal or monacid salt by the addition of phos¬ 
phoric or hydrochloric acid, and evaporating down. It then 
crystallizes out in colourless tablets, which, with a small quantity of 
water, give a crystalline precipitate of the monacid phosphate, but 
which dissolve in a larger quantity of the solvent. It is the chief 
constituent of the manure called superphosphate of lime, or simply 
superphosphate. 


Calcium Carbonate: C 0 * 0 2 Ca.—This compound is very 
widely distributed on the surface of the earth, and in very different 
forms. It is prepared artificially by precipitating a solution of a 
calcium salt with sodium carbonate, and then forms a white 
amorphous powder, which soon becomes crystalline. It is in¬ 
soluble in pure water, even when hot, but dissolves in water con¬ 
taining carbonic acid with the formation of an acid carbonate. 


Calcium Carbonate. 


399 

Calcium carbonate easily dissolves in dilute hydrochloric, nitric, 
or acetic acid, with evolution of carbonic acid; and when heated 
to redness loses its carbonic acid, and is transformed into calcium 
oxide. 

Calcium carbonate is dimorphous : as calcite it crystallizes in 
the hexagonal system, usually in rhombohedra, while as arragoniie 
it belongs to the rhombic system. Both forms may be produced 
artificially. If calcium carbonate is precipitated from a solution 
at the ordinary temperature, the precipitate is at first amorphous, 
but soon becomes crystalline, and then consists of minute rhom¬ 
bohedra like calcite; the same crystals are also formed when the 
clear solution of calcium carbonate in carbonic acid is allowed to 
evaporate slowly. But if the calcium carbonate is precipitated 
from a boiling-hot solution instead of from a cold one, the crys¬ 
talline precipitate then consists of small rhombic prisms, resembling 
those of arragonite. The calcium carbonate deposited from hot 
springs (e.g. Carlsbad) usually has this form. 

The numerous forms in which calcium carbonate occurs in 
nature are designated by different mineralogical names. 

Calcite or calc-spar is nearly pure calcium carbonate crystallized 
in rhombohedra or in some other hemihedral form belonging to the 
hexagonal system. A variety of calcite, which is distinguished by 
its purity and transparency, and which is largely used in optical 
instruments to produce double refraction, is called Iccland-spar. 

Arragonite is the second form in which calcium carbonate 
crystallizes. It is usually found in rhombic prisms, and is consider¬ 
ably rarer than calcite. 

Marble is a granular, crystalline, compact form, produced 
from limestone by the action of underground heat and pressure. 
White marble often looks like loaf-sugar, and consists of small 
crystals of calcite. Marble is found in various colours from white 
to black. 

Cojiipact limestone , of various colours according to the impurities 
which it contains, forms large mountain masses. It has nearly all 
been produced by the agency of marine animals, whose remains are 
often found petrified in it. Oolitic limestone , freestone , consists of 
numerous small globular masses of calcium carbonate cemented 
together, and is so called because it resembles the roe of a fish. 
Lithographic stone is a compact, fine-grained form of limestone, 
found near Solenhofen, in Bavaria, and largely used in the processes 
of lithography. 


400 Text-Book of Inorganic Chemistry. 

Chalk is a soft, earthy variety of calcium carbonate, and from 
its white colour is apparently very pure. It consists chiefly of the 
minute cases of various species of foraminifera, and has been formed 
in deep seas. That it still contains organic traces of these animals 
is readily proved by dissolving it in dilute hydrochloric acid. The 
carbonic acid evolved possesses an unpleasant odour, and a brown 
organic residue remains behind. 

Calcareous tufa or travertine is a porous form of calcium car¬ 
bonate, produced by the gradual evaporation of a carbonic acid 
solution of the carbonate. Of a similar origin are the stalagtites 
found hanging from the roofs of limestone caves, and other similar 
deposits. 

Finally, calcium carbonate is a constituent of many products of 
animal life— e.g. of egg-shells, and of the shells of oysters, and 
other mollusca. On some parts of the coast, where lime is scarce, 
oyster-shells are burnt to obtain quick-lime. 

The technical applications of calcium carbonate are exceedingly 
numerous and important. Marble is a valuable material for the 
builder and sculptor. Chalk is used for writing, and, when finely 
powdered, as a pigment. Compact limestones are employed for 
the manufacture of glass, and as a flux to produce a fusible slag in 
the extraction of many metals from their ores. And, finally, im¬ 
mense quantities of limestone are used for the manufacture of 
quick-lime and then of mortar. 

Bleaching- Powder. (Chloride of Limei) — The chief compound 
contained in this important substance has (p. 117) the composition : 

Ca(OCl)Cl or Caj^^ but the pure compound corresponding to 

this composition has not yet been prepared. Bleaching powder 
is a soft, white, amorphous powder, with a faint odour of hypo- 
chlorous acid, soluble in water with decomposition into calcium 
chloride and calcium hypochlorite and easily decomposed by all 
acids. 

Bleaching powder is manufactured on a large scale by saturat¬ 
ing dry slaked-lime with chlorine. The gas is generated in leaden 
retorts and led into chambers containing shelves on which the 
slaked-lime is placed in thin layers. As soon as the lime ceases 
to absorb the gas, the process is finished ; the chambers are then 
opened and the bleaching powder preserved in well-closed vessels. 


Bleaching Powder. 401 

A solution of bleaching powder in water does not itself bleach, 
or only very slowly as it is decomposed by the carbonic acid of the 
air. If, however, a dilute acid is added, the bleaching-powder is 
decomposed, yielding hypochlorous acid, and then free chlorine. 
The articles to be bleached are first thoroughly cleansed, dipped 
into a dilute solution of bleaching powder, and then into a very dilute 
solution of hydrochloric acid, which, with the bleaching powder 
they contain, liberates free chlorine (p. 117). Thorough washing is 
necessary afterwards, as any small traces of chlorine which remain 
behind gradually rot the stuff. But as it is impossible to completely 
remove the last traces of chlorine, it is customary to dip the articles 
in a solution of some substance which acts chemically upon it— 
called an antichlor. Such a substance is sodium thiosulphate 
(P- 370). 

The formula given above for bleaching powder and the assump¬ 
tion that it is decomposed by water are supported by many facts. 
The action of this substance on bodies which easily take up chlorine 
and oxygen is quite different in the absence and in the presence 
of water. Absolute alcohol and dry bleaching powder react ener¬ 
getically upon one another and produce first ethyl hypochlorite, 
and then acetic ether, but if aqueous alcohol is used, the- chief 
product is chloroform. 

We have already seen (p. 117) that a solution of bleaching 
powder yields calcium chloride and chlorate when boiled. 

Calcium Silicate : SiO-0 2 Ca.—The normal salt of this com¬ 
position is produced as a gelatinous precipitate, drying to a white 
amorphous powder, when calcium chloride in solution is added to 
a solution of sodium silicate. This compound is found in nature 
as the mineral wollastonite , and united with other silicates as 
double salts. To these belong many of the zeolites , lime-garnet , 
datolite , &c. Similar double silicates, but not crystalline, are the 
chief varieties of 


GLASS. 

This important substance is essentially a double silicate of potash 
or soda and lime, in which the lime is sometimes displaced by lead 
oxide. It is distinguished by its transparency, hardness, fusibility 
and resistance to many solvents which attack other substances, 
even the metals, and is therefore indispensable to the chemist. 

D D 


402 


Text-Book of Inorganic Chemistry. 

Glass is manufactured by fusing together a mixture of calcium 
carbonate, potassium (or sodium) carbonate and silica at a high 
temperature, in specially constructed furnaces and crucibles. At 
a bright-red heat the silica expels the carbonic acid from the 
carbonates, producing a double silicate of calcium and potassium 
or sodium, which on continued heating forms a homogeneous liquid 
and on cooling solidifies to a transparent glass without a trace of 
crystalline structure. The purer the materials employed, the purer 
and less coloured is the glass obtained. In order to give the 
required shape to the glass it is partly cast and partly blown. 
Large slabs of thick glass (plate-glass) are obtained by pouring 
the liquid substance on to horizontal slabs of iron, and the com¬ 
moner kinds of glass vessels (bottles, tumblers, &c.), especially the 
thicker ones, are cast in moulds, just as liquid iron is cast. The 
majority of glass vessels are, however, made of what is called 
blown glass. The workman takes a mass of semi-fluid glass on 
the end of an iron tube, and by blowing, combined with various 
movements which cannot be well described, gives to it almost any 
desired shape. Most of the various shaped vessels used in daily 
life, and in the laboratory—wine-glasses, bottles, funnels, flasks, 
&c.—are made in this way. 

Glass is a brittle substance, and if of considerable thickness, or 
rapidly cooled, easily breaks when subjected to rapid changes in 
temperature. In order to allow the molecules of the glass to arrange 
themselves regularly, that the tension on cooling may be the same 
throughout the entire thickness, it must be annealed, or allowed to 
cool slowly. For this purpose, the glass vessels, after they have 
been made, are brought into the annealing furnaces, where the glass 
is heated nearly to the temperature at which it becomes soft, and 
then allowed to cool very slowly during several days or even 
weeks. The influence which rapid cooling has on the properties 
of the glass is well seen in what are called Rupert’s drops, which 
are made by dropping liquid glass into cold water. Owing to the 
rapid cooling, they are hard, and stand pressure or even a blow. 
But if slightly scratched, or if a small piece is broken off the thin 
end, the whole drop breaks up into small pieces of glass. 

A few years ago a variety of quickly annealed glass was intro¬ 
duced into trade, and highly recommended for various purposes. 
This toughened glass, as it is called, is harder than common glass, 
and bears rapid" changes of temperature better. It is prepared by 
dipping the freshly blown hot glass into a bath of hot oil or paraffin, 


Glass. 


403 

a process which resembles that by which Rupert’s drops are formed, 
except that the cooling is less sudden. Vessels prepared in this 
way may be dropped from a considerable height on to the floor 
without breaking, and a glass dish filled with cold water may be 
placed on glowing charcoal, but, like Rupert’s drops, they will not 
stand the slightest scratch. A sharp grain of sand is often 
sufficient to cause their violent fracture into a number of small 
pieces, and it sometimes happens that vessels of toughened glass, 
which have been used for years, and have withstood considerable 
changes of temperature without damage, suddenly break up into a 
thousand pieces. Toughened glass vessels, from which so much 
was expected at the time of their introduction, should therefore 
never be used, especially for chemical apparatus. 

Various kinds of glass are distinguished from one another 
according to their constituents, and according to the purity of the 
substances employed in their manufacture. Potash , or hard-glass , 
is a double silicate of calcium and potassium, and is largely manu¬ 
factured in Bohemia. It is less acted on by chemical reagents than 
any other kind of glass, and when the proportion of silica is high, 
is difficultly fusible. The glass tubes used for combustions in 
organic analysis are made from this difficultly fusible glass. 
Crown , or soda-glass , contains soda instead of potash, and is the 
common window-glass. It is more easily fusible than potash-glass, 
and more readily attacked by acids. Bottle-glass is a coarse form 
of crown-glass, made with impure materials, and more or less 
coloured by the presence of iron. It also contains not inconsider¬ 
able quantities of aluminium silicate. Flint-glass is a double 
silicate of lead and potassium— i.e. it is potash-glass in which the 
calcium has been displaced by lead. It is prepared by fusing 
together silica, potassium carbonate, and lead oxide (litharge), and 
is distinguished by its lustre and high index of refraction. Flint- 
glass is the most easily fusible of all glasses, and the most easily 
attacked by acids ; large quantities of it are manufactured in 
England and France. Flint-glass is especially valuable for optical 
purposes, owing to its high index of refraction. The refractive 
power of the glass can be increased by using a larger proportion 
of lead oxide ; such glass, called strass , is used for the manufacture 
of artificial gems. 

Finally, the various slags obtained in metallurgical operations 
are also a species of glass. In smelting some of the heavy metals, 
especially iron, calcium carbonate is often added as a flux to form 

d d 2 


404 Text-Book of Inorganic Chemistry . 

a fusible slag, which removes the earthy impurities, swims on the 
surface of the heavier metal, and so protects it from the oxygen of 
the air. The slags consist of the silicates of the alkalies, of lime, 
alumina, magnesia, iron, &c., and when cold yield a strongly 

coloured, often opaque, vitreous mass. 

Glasses of various colours are produced by the addition ot 
different metallic oxides to the liquid glass. The silica of the glass 
unites, like borax, with these oxides, forming coloured silicates 
Blue glass is produced by the addition of small quantities of 
cobalt; green by uranium or chromium; olive-green by ferrous 
oxide ; yellow or brown by ferric oxide ; amethyst by manganous 
oxide ; red by cuprous oxide, or a trace of gold, &c. Coloured 
glass often consists of colourless glass with the coloured glass on 
one side only. Such glass is said to be flashed . White, opaque, 
or milk- glass, is made by adding some white infusible powder, such 
as bone-ash, or cryolite, to the molten glass. 

Calcium Oxalate-. 0 2 Ca+H 2 0 , is contained in many 

plants, especially in the sorrels, and is obtained by precipitating a 
solution of calcium chloride with ammonium oxalate. The preci¬ 
pitate is crystalline, and loses its water of crystallization at ioo°. 
The salt is insoluble both in water and in dilute acetic acid, and 
serves, therefore, for the estimation of calcium and of oxalic acid. 

calcium Chloride : CaCl 2 .—This salt, which is distinguished 
by the energy with which it attracts water, and which is, therefore, 
largely used for drying purposes in chemistry, is obtained as a bye- 
product in many operations—^, in the preparation of ammonia 
from ammonium chloride and quick-lime (p. 184)5 an d °f carbonic 
acid from calcium carbonate and hydrochloric acid. The acid 
liquid resulting from the latter process when evaporated down to a 
small bulk deposits large rhombic crystals of the composition : 
CaCl + 6 H 2 0 . These crystals deliquesce in the air and dissolve 
readily in water with a considerable fall in temperature; and 
when mixed with snow or powdered ice the temperature may fall 
as low as -45°. If this crystalline compound is heated, it melts 
in its water of crystallization and loses 4 molecules of water, leaving 
a white porous mass, which readily absorbs water again. At 
a red heat the salt loses all its water and fuses to a clear liquid. 
This fused mass is poured on to a clean iron plate, broken into 


Calcium Chloride . 


405 


pieces when cold, and so brought into trade. Fused calcium 
chloride dissolves in water with evolution of heat ; it is also solu¬ 
ble in alcohol. It energetically absorbs dry ammonia gas, pro¬ 
ducing a compound : CaCl 2 + 8NH 3 , which is decomposed into its 
components either when heated or when brought into contact with* 
water. Calcium chloride cannot, therefore, be used to dry gaseous 
ammonia. 

Calcium Fluoride : CaF 2 , is the mineral fluor-spar . It crystal¬ 
lizes in the regular system, but is also found in compact masses. 
When finely powdered it is used for the preparation of hydrofluoric 
acid, and as a flux to produce a fusible slag—hence its name (from 
fluo , I flow). Calcium fluoride is quite insoluble in water, and is 
therefore precipitated when a solution of a calcium salt is mixed 
with one of a soluble fluoride. 

Calcium Sulphide : CaS, is obtained as a yellowish-white earthy 
mass by glowing quick-lime in a stream of sulphuretted hydrogen, 
or, less pure, by igniting a mixture of powdered calcium sulphate 
and charcoal, when calcium sulphide and carbonic oxide are 
rmed. The compound so produced is insoluble in water, but 
gradually decomposes in contact with this liquid into calcium 
hydrate and sulphydrate :— 

2 CaS + 2H 2 0 = Ca(OH) 2 + Ca(SH) 2 . 

Calcium sulphide, like the other sulphides of the alkaline earths, 
shines in the dark after it has been exposed to a bright light—in 
other words, it phosphoresces. It is the chief constituent of Bal¬ 
main’s luminous paint, in which form it is used to render match¬ 
boxes, buoys, and other objects luminous in the dark. 

Calcium Sulphydrate : Ca(SH) 2 , is easily obtained in solution 
by saturating milk of lime with sulphuretted hydrogen. It possesses 
the property of softening hair and changing it into a gelatinous 
mass. Mixed with other substances, it is used by some Oriental 
nations as a depilatory. 

Calcium Pentasulphide : CaS 5 .—A mixture of this salt and 
calcium thiosulphate may be prepared by boiling milk of lime with 
an excess of flowers of sulphur, and filtering off the reddish-coloured 
liquid. Like potassium pentasulphide, hydrochloric acid decom¬ 
poses it with a copious precipitate of milk of sulphur (p. 145). 


406 


Text-Book of Inorganic Chemistry. 


Detection of Calcium Compounds. 

Calcium compounds when introduced into a non-luminous gas- 
flame give a reddish colour, somewhat yellower than that produced 
by the compounds of strontium. The spectra of the two flames are, 
however, so different, that they may be very easily distinguished 
from one another. If the compound is an insoluble one, it must 
be first moistened with hydrochloric acid before it is introduced 
into the flame. 

Calcium salts closely resemble those of strontium and barium. 
The sulphates are all difficultly soluble in water, that of calcium 
being the most soluble. The carbonates are all equally insoluble 
in water, whence it follows that they are all precipitated together by 
sodium or ammonium carbonate. The oxalates are also all insoluble 
in water, but that of calcium is distinguished from the other two 
by its insolubility in acetic acid. Of the anhydrous nitrates only 
that of calcium is soluble in absolute alcohol, a fact which is utilized 
to separate calcium from strontium and barium. 


STRONTIUM. 

Chemical Symbol : Sr .—Atomic Weight : 87*5. 

The compounds of this element closely resemble those of 
calcium and barium, especially the latter, and strontium with its 
compounds forms a natural transition from calcium to barium. 

The most important minerals containing strontium are celestine 
or strontium sulphate, and strontianite , strontium carbonate. The 
latter mineral is called from Strontian in Scotland, where it was 
first found in quantity, hence the name given to the element. 

Recently, large deposits of strontianite have been found in 
Westphalia, and this mineral is now almost exclusively used for 
he preparation of strontium compounds. 

The metal may be obtained by the electrolysis of the fused 
chloride. Its specific gravity is 2*5, it is harder than lead, and is 
malleable. Strontium, like calcium, is a yellow metal and melts at 
a red heat. 

strontium Oxide : SrO, cannot be prepared, like calcium 
oxide, by glowing the carbonate, as this compound is only slowly 



Strontium. 


407 


and incompletely decomposed by heat. It is obtained, like barium 
oxide, by heating its nitrate to redness, and forms a grey, porous 
mass. 

This oxide unites with water and produces the hydrate: 
Sr(OH) 2 , which crystallizes from water as the compound Sr(OH) 2 + 
8 H 2 0 . Strontium hydrate is more soluble in water than the 
calcium compound, but less than that of barium ; 1 part dissolves in 
about 60 parts of cold water. The solution is strongly alkaline. 

Stro?itium Peroxide : Sr 0 2 is obtained, like the calcium com¬ 
pound, by precipitating a solution of strontium hydrate with 
hydrogen peroxide. Strontium oxide does not unite directly with 
oxygen. 


strontium Sulphate : S 0 2 * 0 2 Sr, is found in nature, as the 
mineral celestine , in transparent, rhombic prisms, often of consider¬ 
able beauty. The salt may be obtained as a white powder by 
precipitating a solution of a strontium salt with sulphuric acid or 
a soluble sulphate. It is slightly soluble in water, but less so than 
calcium sulphate ; 1 part of the salt requires about 6,000 parts of 
water for solution. 


Strontium Nitrate 


NQ 2 } * s distinguished from calcium 

nitrate by its insolubility in alcohol, and in other respects resembles 
the barium salt. It is easily obtained by dissolving strontianite in 
dilute nitric acid and evaporating down. As strontium compounds 
impart a red colour to flames into which they are introduced, its 
nitrate is used instead of potassium nitrate to produce an inflam¬ 
mable powder which burns with a red light. 


Strontium Carbonate : CO • 0 2 Sr, is obtained artificially like the 
calcium compound, which it closely resembles. It occurs in nature 
as the mineral strontianite ; either in rhombic prisms or in indistinct 
crystalline masses. 


Strontium Chloride : SrCl 2 + 6 H 2 0 , crystallizes in needles, is 
easily soluble in water and slightly deliquesces in the air. Like 
calcium chloride, it is soluble in strong alcohol. 


Strontium Sulphide : SrS, is obtained from celestine by heating 
a mixture of the powdered mineral with charcoal, and forms the 
starting-point in the preparation of strontium compounds from this 
source. Like calcium sulphide, it is luminous in the dark. 


408 


Text-Book of Inorganic Chemistry. 


Detection of Strontium Compounds. 

Strontium compounds impart an intense crimson colour to the 
non-luminous flame, which can at once be distinguished from the 
lithium or calcium flame by the spectroscope. The spectrum of 
strontium is especially characterized by a fine line in the bright 
blue. 

Strontium salts are distinguished from those of calcium by the 
fact that calcium sulphate produces a precipitate of strontium 
sulphate on standing. They give no precipitate with potassium 
chromate, especially if the solution contains free acetic acid, 
although the chromate is less soluble than the calcium salt. 

Strontium nitrate is insoluble in strong alcohol and thus differs 
from calcium nitrate. This property affords a ready means of 
quantitatively separating the two salts. 


BARIUM. 

Chemical Symbol : Ba .—Atomic Weight : 137. 

This element, which derives its name from the high specific 
gravity of its compounds (@apvs = heavy), is found in nature like 
strontium chiefly as the sulphate, barite or heavy s£ar, and as the 
carbonate, witherite. The former compound is the source from 
which nearly all barium compounds are prepared. 

Metallic barium is exceedingly difficult to prepare. It is said 
to have been obtained by electrolysis of the fused chloride, and is 
described as a yellow metal, melting at a bright red heat and de¬ 
composing water at the ordinary temperature. 

Barium, like calcium and strontium, is a dyad element in nearly 
all its compounds ; in the peroxide : Ba 0 2 , it is a tetrad. 

Barium Oxide (Baryta) : BaO, cannot be prepared by heating 
the carbonate, as the carbonic acid is more firmly united even than 
in strontium carbonate, and is only very slightly expelled even at 
the highest temperatures. Barium oxide is best prepared by heat¬ 
ing barium nitrate. A Hessian crucible is about half filled with 
the salt and heated first gently to expel most of the nitric peroxide, 
and then, when the frothing has subsided, to bright redness. 
The oxide which remains behind is a grey porous mass, which has 



Barium . 


409 


apparently been fused, but is really infusible at this temperature. 
Occasionally it has a greenish colour, due to manganese derived 
from the crucible which has been converted into barium manga- 
nate. When mixed with water, it evolves a considerable amount 
of heat and is converted into barium hydrate. When gently 
heated in the air or oxygen, it absorbs considerable quantities of 
oxygen and is changed into barium peroxide. 

Barium Hydrate : Ba(OH) 2 —The preparation of barium 
hydrate from barium oxide is tedious and unsuitable if large 
quantities are required. It is better to employ the sulphide, 
which is easily obtained from the sulphate by heating with 
powdered charcoal. For this purpose a solution of the sulphide 
in water is boiled with copper oxide until a portion of the liquid 
no longer blackens a solution of lead acetate. The clear solution 
is then filtered off from the insoluble copper sulphide. The 
reaction is expressed in the following equation :— 

BaS + CuO + H 2 0 = Ba(OH) 2 + CuS. 

From the hot solution so obtained, crystals of the compound : 
Ba( 0 H) 2 + 8 H 2 0 separate out on cooling in tablets or prisms. 
Barium hydrate may also be prepared from the sulphide by first 
converting this salt into the carbonate by moist carbonic acid- 
sulphuretted hydrogen being set free—and then acting on the car¬ 
bonate with superheated steam— i.e. steam at high pressure. The 
carbonic acid then passes away with the steam, and barium hydrate 
remains behind. The crystalline compound parts with its water 
of crystallization when heated, but, unlike calcium hydrate, cannot 
be converted into the oxide by heat. 

Barium hydrate dissolves in about 20 parts of cold water, but 
requires less than 1 part of hot water for solution. It is thus much 
more easily soluble than calcium or strontium hydrate. The solu¬ 
tion ( baryta-water ) is strongly alkaline and precipitates the weaker 
bases from their solutions like caustic potash. When exposed to 
the carbonic acid of the air it becomes turbid more rapidly than 

I lime-water. Baryta is a weaker base than potash or soda, and is 
therefore precipitated by these alkalies, not by ammonia, from 
concentrated solutions of its salts. 

Barium oxide possesses the remarkable property of forming an 
insoluble compound with cane-sugar; and on this account large 
quantities are used by sugar refiners. Strontium oxide also pos¬ 
sesses this property. 



410 Text-Book of Inorganic Chemistry. 

Barium Peroxide : Ba 0 2 .—Barium oxide, when heated iofaint 
redness in a stream of air or oxygen, increases in weight owing 
to absorption of oxygen, and becomes gradually converted into 
barium peroxide, which cannot be distinguished by its external 
properties from the oxide itself. At a slightly higher temperature 
than that required for its production, but still far below bright red¬ 
ness, it again parts with an atom of oxygen, and is reconverted 
into the oxide, which can again unite with oxygen, and so on. 
Experiments made on a large scale to ascertain whether in this 
way large quantities of pure oxygen might not be obtained from 
the air have not been altogether successful. It appears that the 
oftener the oxide is converted into peroxide, the less becomes 
the energy with which it unites with oxygen, until it finally loses 
this property altogether. The cause of this is possibly because 
the oxide gradually becomes more compact, and so exposes less 
surface to the action of the gas. Recently it has been discovered 
that barium peroxide parts with its oxygen at the same temperature 
at which it is produced if the pressure is reduced. Or, in other 
words, to obtain a supply of pure oxygen, the barium oxide is 
kept at a constant temperature, and the pressure only varied. 
Possibly some use may be made of this interesting fact. 

Barium peroxide prepared as above described is never pure, 
but always contains more or less barium oxide. It may be 
purified in the following way. It has already been stated (p. 90) 
that barium peroxide is distinguished from manganese and other 
peroxides by the fact that it gives hydrogen peroxide, not chlorine, 
when treated with hydrochloric acid, and its purification depends 
upon this reaction. The impure barium peroxide, obtained by 
heating barium oxide in a stream of air, is finely powdered and 
rubbed up with water in a mortar, with which it unites chemically. 
The mixture is next gradually added to dilute hydrochloric acid, 
until the liquid only reacts faintly acid, and the clear solution, con¬ 
taining barium chloride and hydrogen peroxide besides the excess 
of acid, filtered off. On adding an excess of baryta-water to this 
liquid, the pure hydrated peroxide separates out in lustrous tablets 
of the composition : Ba 0 2 + 8 H 2 0 . These crystals when washed, 
and dried between filter paper and over sulphuric acid, gradually 
lose all their water and leave pure barium peroxide behind. The 
pure compound remains unchanged in the air, gives up one-half of 
its oxygen when heated, and unites with water to form the hydrated 
compound. 


Compounds op Barium 


4i 1 


Barium Sulphate : S 0 2 * 0 2 Ba.—This salt, the most impor¬ 
tant naturally occurring compound of barium, is usually found 
in compact masses, and sometimes in tabular crystals belonging 
to the rhombic system. It has a specific gravity of 4*4, and is 
known as barite or heavy-spar. It may be readily prepared arti¬ 
ficially, as a heavy, white precipitate, by adding sulphuric acid or 
a sulphate to a soluble barium salt. Barium sulphate is quite 
insoluble both in water and dilute acids, and is therefore used to 
precipitate barium or sulphuric acid from a liquid. Concentrated 
sulphuric acid dissolves it to some extent, but on dilution with 
water it is again precipitated, like lead sulphate. Precipitated 
barium sulphate is used as a white pigment under the name of 
permanent white ; it is also largely used to adulterate white-lead, 
for which purpose its high specific gravity, combined with its 
cheapness, makes it admirably adapted. 



Barium Nitrate : 


erite or barium sulphide in nitric acid, and is deposited from the 
hot solution in brilliant, heavy octahedra, without water of crystal¬ 
lization. The salt is tolerably easily soluble in water, especially 
when hot, but less soluble in the presence of free acid ; hence, on 
adding nitric acid to an aqueous solution, a precipitate is produced. 
It is insoluble in alcohol. Barium nitrate is poisonous, like nearly 
all the barium salts. 

Barium Carbonate : CO • 0 2 Ba, occurs in nature as the mineral 
witherite , and is thrown down as a white precipitate on adding 
sodium carbonate to a solution of a barium salt. It is insoluble in 
water, but slightly soluble in the presence of carbonic acid ; and, 
notwithstanding its insolubility, is poisonous. 

Barium Chloride : BaCl 2 + 2H 2 0. — This compound is one of 
the least soluble of the barium salts. It is prepared in a similar 
manner to the nitrate, and crystallizes in rhombic plates. At ioo° 
it loses its water of crystallization and melts at a red heat. Alcohol 
does not dissolve it—a reaction which is employed to separate 
barium from calcium and strontium. Its aqueous solution is largely 
used in analytical chemistry as a test for various acids. 


412 Text-Book of Inorganic Chemistry. 

Barium Sulphide : BaS ; and Barium Sulphydrate : Ba(SH) 3 . 
Barium sulphide is obtained as a white amorphous mass when 
barium oxide is heated in a stream of sulphuretted hydrogen, or, 
less pure, by reducing the sulphate with charcoal. An intimate 
mixture of finely powdered heavy-spar and charcoal is mixed with 
starch paste and formed into balls. These are then dried in the air, 
placed between alternate layers of charcoal in a furnace, the bottom 
layer of which is ignited. When cold, they are extracted and 
digested with water. The solution, especially if prepared hot, does 
not contain barium sulphate, but a mixture of barium sulphydrate 
and hydrate, which crystallize together :— 

2 BaS + 2H 3 0 = Ba(SH)» + Ba(OH) 2 . 

If sulphuretted hydrogen is led into this solution, the hydrate is 
converted into sulphydrate, and this salt alone crystallizes out on 
evaporation. As heavy-spar is a much commoner mineral than 
witherite, barium sulphide is the compound from which nearly all 
6 ther barium salts are manufactured. 

Detection of Barium Compounds. 

Compounds of barium when introduced into a Bunsen flame 
produce a yellowish-green colour, the spectrum of which is more 
complicated than that of either calcium or strontium, and is dis¬ 
tinguished by several lines in the green. 

In their chemical characters the barium compounds closely 
resemble those of calcium and strontium, but it is still easy to dis¬ 
tinguish them from one another. Barium sulphate is so insoluble 
in water, that on adding a solution of calcium sulphate to one of a 
barium salt an immediate turbidity is produced. In the same 
manner calcium and strontium chromates are much more soluble 
than barium chromate, and potassium chromate added to a dilute 
solution of barium chloride produces a pale yellow precipitate of 
barium chromate, but no precipitate is produced with calcium (or 
strontium) chloride. Barium can be separated from calcium or 
strontium by digesting the perfectly dry chlorides with absolute 
alcohol. Calcium and strontium chlorides then dissolve while 
barium chloride remains undissolved. 



Magnesin m. 413 

MAGNESIUM. 

Chemical Symbol : Mg.— Atomic Weight : 24. 

The compounds of magnesium are not less widely distributed 
than those of calcium, but are not found in such large quantities. 
Among the more important minerals containing magnesium are : 
magnesite (magnesium carbonate: CO• 0 2 Mg), dolomite (mag¬ 
nesium-calcium carbonate : CO • 0 2 Mg + CO • 0 2 Ca), talc, steatite or 
soap-stone, serpentine, and meerschaum (all essentially magnesium 
silicate). Other silicates which are rich in magnesium are : augite, 
hornblende, asbestos, olivine, and biotite or magnesium mica. Mag¬ 
nesium also occurs in combination with sulphuric acid as Epsom 
salt in many mineral waters and in the sea. 

Magnesium is a malleable silver-white metal of specific gravity 
= 175. It melts at a low red-heat, and at a higher temperature 
can be distilled. It remains almost unchanged when exposed to 
the air, and only becomes covered with a thin, grey coating of the 
oxide which preserves the metal from further oxidation. When 
heated in the air above its melting-point it catches fire and burns 
with a brilliant white light, producing its only oxide—magnesia. 
A wire or thin ribbon of magnesium very easily takes fire and 
burns when one end is held in the flame of a Bunsen burner. The 
light of burning magnesium is so intense and so rich in chemical 
rays that a mixture of chlorine and hydrogen explodes when ex¬ 
posed to it, just as it does in bright sunlight. 

Magnesium does not decompose water at the ordinary tempe¬ 
rature, or at all events only very slowly ; at about 30° a feeble 
evolution of hydrogen begins, which becomes more rapid as the 
temperature rises. Dilute acids, as well as ammonium chloride or 
carbonate, readily dissolve it. 

Magnesium may be prepared either by the electrolysis of the fused 
chloride, or else by decomposing the chloride with sodium. The 
electrolysis of the chloride may be performed in a porcelain crucible 
partially divided into two divisions by a piece of unglazed porcelain 
(fig. 63). When the crucible has been filled with the fused chloride, 
the two electrodes, which are made of gas-carbon and fitted in the 
lid (fig. 64), are introduced and connected with the poles of a 
powerful battery (about 8 to 10 Bunsen cells). The metal which 
is then produced is lighter than the fused chloride, and would rise 
to the surface and burn in the air if the negative pole were not 


4 i 4 Text-Book of Inorganic Chemistry. 

provided with a series of notches which catch the globules of metal 
in their ascent. 

For the reduction of magnesium from the chloride by sodium, 
a mixture is prepared of six parts of anhydrous magnesium chloride, 



Fig. 63. Fig. 64. 


one part of fluor-spar, one part of fused potassium-sodium chloride, 
and one part of sodium cut into small pieces, and the whole thrown 
into a red-hot Hessian crucible. As soon as the first violent reac¬ 
tion is over, the closed crucible is removed from the furnace, and 
when it is no longer red hot, its contents are stirred with a piece of 
tobacco-pipe to unite the smaller globules of metal into larger ones. 
When the crucible is cold it is broken open, and the globules of 
metal picked out and washed with water. 

Magnesium is a dyad in all its compounds ; it unites in only 
one proportion with oxygen. 

Magnesium Oxide. Magnesia : MgO. 

This, the only oxide of magnesium, is a white, light powder, 
which is infusible at the highest temperatures of our furnaces. 
When strongly heated it becomes much denser, and its specific 
gravity increases from 3*2 to 3*6. It is especially distinguished 
from the oxides of calcium, strontium, and barium by its almost 
complete insolubility in water. But although one part of magnesia 
requires more than 50,000 parts of water for solution, it still pro¬ 
duces a distinctly blue spot when moistened with water and placed 
on red litmus paper. It unites with water with a very slight evolu¬ 
tion of heat, forming the hydrate, from which the water is again 
easily expelled on heating. It attracts carbonic acid from the air 
(when in the pulverulent state), and, being a strong base, forms 
compounds with all acids. 



4 T 5 


Magnesium Sulphate. 

Magnesium oxide is obtained by gently glowing the basic car¬ 
bonate, known to pharmacists as magnesia alba , and then gives 
a fine white powder called magnesia usta or calcined magnesia . 
The oxide is produced as a solid, compact mass on glowing the 
compact natural form of magnesium carbonate, magnesite. Mag¬ 
nesium oxide gradually hardens in contact with water, and forms 
a compact mass of the hydrate. 

Magnesium Hydrate : Mg(OH) 2 , is precipitated on adding 
caustic soda or ammonia to a solution of magnesium sulphate ; 
when washed and dried, it yields a white powder. It may be 
heated to ioo° without losing its water of hydration, and slowly 
attracts carbonic acid from the air. It is found in nature as the 
mineral brucite. 

Magnesium Sulphate. Epsom Salt : S 0 2 * 0 2 Mg + 7 H 2 0 .— 
This salt differs from the sulphates of calcium, strontium, and 
barium, by its easy solubility in water, and by the readiness with 
which it crystallizes. Considerable quantities of magnesium sul¬ 
phate united with one molecule of water are found in the Stassfurt 
beds, and are known as Kieserite : S 0 2 * 0 2 Mg + H 2 0 . 

Magnesium sulphate is further contained in sea-water and in 
many highly valued mineral waters, among which may be men¬ 
tioned those of Epsom near London, Seidlitz, Friedrichshall in 
Germany, and Hunyadi Jdnos in Hungary. The salt may be 
obtained artificially by decomposing magnesite or dolomite with 
dilute sulphuric acid, and is a bye-product in manufactories of 
aerated waters when either of these substances is used to prepare 
carbonic acid. If dolomite is used, calcium sulphate remains un¬ 
dissolved. 

Magnesium sulphate crystallizes in colourless rhombic prisms, 
and possesses an unpleasant saline and bitter taste. It is easily 
soluble in water, insoluble in alcohol, and when heated melts in its 
water of crystallization, of which it loses six molecules under 150°. 
The last molecule is only expelled at a temperature above 200°. 
This molecule of water—so-called water of constitution —can be 
displaced by salts, producing double compounds, containing six 
molecules of water. Thus, when solutions of magnesium and 
potassium sulphates are mixed, a double salt of the composition : 
S0 2 -0 2 Mg +S 0 2 ( 0 K) 2 + 6 H 2 0 , crystallizes out. In this respect 
magnesium sulphate is closely allied to other sulphates of dyad 


416 Text-Book of Inorganic Chemistry. 


metals which also crystallize with seven molecules of water (the 
vitriols), and thus connects the metals of the alkaline earths with 
the heavy metals. 

Owing to its well-known action as a mild and safe purgative, 
magnesium sulphate is a valuable and much used medicine, and 
mineral waters containing it are among those of which the largest 
quantities are consumed. 


Magnesium Phosphates. 


The normal salt : (PO) 2 O e Mg 3 , is obtained as a white precipi¬ 
tate on mixing a solution of normal sodium phosphate w r ith one of 
magnesium sulphate. If common (monacid) sodium phosphate is 

added to magnesium sulphate the monacid phosphate : Poj OH b + 

7 H o 0 , is produced, which separates from'concentrated solutions as 
a white, amorphous powder, and is slowly deposited from dilute 
solutions in the crystalline form. The salt is difficultly soluble 
in water, but easily dissolves in acids and in salts of ammonium. If 
ammonia is added to a solution of magnesium phosphate in am¬ 
monium chloride, or to a mixture of magnesium sulphate, am¬ 
monium chloride, and sodium phosphate, a crystalline precipitate 
at once separates which is insoluble both in ammonium salts and 
in water containing free ammonia. This compound is a double 
phosphate of magnesium and ammonium and has the composition : 

PO 1 + 6PLO. If heated it loses its water of crystalliza- 

^ONH 4 

tion and ammonia, while magnesium pyrophosphate remains be¬ 
hind, according to the equation :— 




+ 


The insolubility of this ammonium-magnesium phosphate renders 
it a suitable compound for quantitatively estimating both magnesium 
and phosphoric acid. The salt is often found in decomposing 
urine, and is sometimes deposited in the bladder as a constituent 
of urinary calculi. It is also contained in guano and occasionally 
in old dung-heaps, often forming large, well-developed crystals 
{struvite). 

Ammonium-Magnesium Arsenate'. AsO + 6 H 2 0 , is 


Magnesium Carbonate. 417 

prepared like the corresponding phosphate by mixing solutions 
of magnesium sulphate, ammonia, and sodium arsenate. Its pro¬ 
perties closely resemble those of the phosphate with which it is 
isomorphous. When heated it also leaves a residue of magnesium 
pyrarsenate. 

Magnesium Carbonate: CO• 0 2 Mg.—This insoluble salt is 
found in nature as magnesite , partly crystallized in rhombohedra 
isomorphous with calcite, and partly as compact crystalline masses. 
Dolomite is a double carbonate of calcium and magnesium car¬ 
bonate. When a magnesium salt is precipitated with sodium car¬ 
bonate, a basic, not the normal salt is produced. This basic 
carbonate, or compound of normal magnesium carbonate and 
magnesium hydrate, has a varying composition, but may be gene¬ 
rally represented either as : 3CO • 0 2 Mg + Mg(OH) 2 or 2CO • 0 2 Mg 
+ M g( OH ) 2 > with varying quantities of water. Acid sodium car¬ 
bonate is produced at the same time in which a small quantity 
of the magnesium carbonate remains dissolved. This white pre¬ 
cipitate is brought into trade as magnesia alba. 

Magnesium carbonate forms double salts with the carbonates of 
the alkalies, which may be obtained by digesting magnesia alba 
with the acid alkaline carbonates. 

Magnesium Borate is deposited on boiling a solution of magne¬ 
sium sulphate with one of borax, but dissolves again on cooling. 
Magnesium borate in combination with magnesium chloride forms 
the mineral boracite , crystallizing in the regular system. 

Magnesium Silicate : Si( 0 2 Mg) 2 .—The normal salt of this com¬ 
position occurs in nature as the mineral olivine. The minerals 
talc , steatite , serpentine , and meerschaum are compounds of mag¬ 
nesium silicate with varying quantities of water. Augite , hornblende, 
and asbestos are double silicates of magnesium and calcium. 

Magnesium Chloride : MgCl 2 .—This salt is contained in sea¬ 
water and many mineral springs. It is obtained pure by dissolving 
magnesia alba in hydrochloric acid, adding chlorine water to oxidize 
the iron, and then digesting the solution with a slight excess of 
magnesia alba to precipitate the iron and alumina. The clear 
solution when evaporated down to a small bulk deposits very de¬ 
liquescent crystals of the composition : MgCl 2 + 6H 2 0. The water 

E E 


4i 8 Text-Book of Inorganic Chemistry. 

cannot be expelled from this hydrated compound by heating since 
the chloride then decomposes into magnesia and hydrochloric 

aCld MgCl 2 + H 2 0 = MgO + 2HCI, 

and a basic chloride is formed. This decomposition can, however, 
be prevented by the addition of ammonium chloride, as a mixture 
of magnesia and ammonium chloride when heated produces mag¬ 
nesium chloride, ammonia and water 

MgO + 2 NH 4 C 1 - MgCl 3 + 2NH3 + HoO. 

The concentrated aqueous solution is therefore mixed with 
enough ammonium chloride to form the double compound : 
MgCl 2 + NH 4 Cl, and then thoroughly dried. The dried mass is 
afterwards gently heated in a Hessian crucible until vapours of 
ammonium chloride are no longer given off and the whole fuses 
quietly. The fused salt is poured on a clean slab of iron, rapidly 
broken into pieces and preserved in well-closed vessels. Some basic 
salt usually remains behind in the crucible. 

Anhydrous magnesium chloride is a white crystalline mass, 
deliquescing in the air. It can be distilled in a stream of hydrogen 
at a bright-red heat, and gives a colourless, soft distillate, which 
solidifies 5 to a crystalline mass of brilliant plates. The salt is easily 
soluble both in water and alcohol. 

Magnesium chloride forms double salts not only with am¬ 
monium chloride, but also with potassium and other chlorides. 
Potassium-magnesium chloride : KC 1 , MgCl 2 , + 6 H 2 0 is found in 
considerable quantities at Stassfurt, and is known as carnallite . 

Detection of Magnesium Compounds. 

Magnesium compounds do not colour a non-luminous gas flame 
when introduced into it, neither do they produce any characteristic 
spectrum. In this respect, magnesium differs from the other 
metals of the alkaline earths, and resembles those of the earths. 
It is further distinguished from the former by the solubility of its 
sulphate, and by the fact that solutions of its salts are not pre¬ 
cipitated by ammonium carbonate, although magnesium carbonate 
is nearly as insoluble in water as the carbonates of calcium, 
strontium, and barium. The reason of this is that all insoluble 
magnesium compounds, except the double phosphate and arsenate 
of ammonium and magnesium, form soluble double compounds 
with ammonium salts. Hence a solution of magnesium sulphate 


Metals of the Earths. 419 

to which ammonium chloride has been added gives no precipitate 
of magnesium hydrate with ammonia. 

From these reactions, magnesium salts may be just as readily 
detected and separated as those of the other alkaline earths. From 
calcium, which commonly accompanies it, it may be separated by 
adding ammonium chloride and ammonia to the mixed solution, 
and then ammonium oxalate or carbonate. All the calcium is 
then precipitated as calcium oxalate or carbonate, and all the 
magnesium is contained in the filtered solution, from which it may 
be thrown down with sodium phosphate as ammonium-magnesium 
phosphate. 

It is remarkable that ammonia (in the absence of ammonium 
salts) partly precipitates magnesia from neutral solutions, while in 
solutions of calcium, strontium, and barium, it produces no pre¬ 
cipitate. 


METALS OF THE EARTHS. 

The chief representative of these metals is aluminium ,* but the 
group also includes a number of metals of less importance, and 
much less known—viz. beryllium , gallium , and indium ; yttrium, 
terbium, erbium and ytterbium ; cerium , lanthanum , and didy- 
miui 7 i; thorium and zirconium. All these metals have a strong 
attraction for oxygen, but considerably less than the metals of the 
alkaline-earths. Their basic oxides are sesquioxides—except those 
of thorium and zirconium. Notwithstanding their strong attrac¬ 
tion for oxygen, they appear to possess only a weak affinity for sul¬ 
phur, and their sulphides (except that of indium) cannot be obtained 
in the wet way. If ammonium or sodium sulphide is added to a 
solution of one of their salts, a precipitate of the hydrate, not the 
sulphide, is produced, and sulphuretted hydrogen is set free :— 
A 1 2 C 1 6 + 3(NH 4 ) 2 S + 6 H 2 0 = 2 A 1 ( 0 H) 3 + 6 NH 4 C 1 + 3 H 2 S. 


ALUMINIUM. 

Chemical Symbol: Al.— Atomic Weight : 27. 

This element is never found free in nature, but combined with 
oxygen and silica it is one of the chief constituents of the super- 

1 The author includes chromium with these metals, but the editor prefers to 
place it after manganese. The editor has also made some slight alterations— 
with additions—in treating of the rarer of these metals. 


E E 2 




420 Text-Book of Inorganic Chemistry. 

ficial layer of the crust of the earth. Felspar, granite, mica, and 
many other rock masses contain aluminium silicate. In combina¬ 
tion with fluorine and sodium fluoride it forms the mineral cryolite. 

Aluminium is a tin-white metal of specific gravity 2-6, tough, 
but malleable; it may be rolled into thin sheets and drawn into 
fine wire. It melts at about 700 °—i.e. at a higher temperature than 
zinc, but lower than silver; it is not magnetic, but conducts an 
electric current better than iron. The metal remains unchanged 
in dry as well as in moist air, becoming covered with an extremely 
thin layer of oxide—even at its melting point it only oxidizes slowly. 
If, however, a large surface is heated in the air— e.g. a piece of 
aluminium foil in the lamp, it burns brilliantly. 

Concentrated nitric acid does not act upon aluminium, but dilute 
sulphuric or hydrochloric acid dissolves it readily with copious 
evolution of hydrogen. It easily dissolves in a solution of caustic 
potash or soda, evolving hydrogen and producing soluble com¬ 
pounds of alumina and potash or soda. The molten alkalies do 
not act upon it, but when fused with potassium carbonate, carbon 
is reduced. 

The strong affinity of aluminium for oxygen and the halogens 
makes its preparation somewhat difficult. Its oxide is not reduced 
when glowed with charcoal, nor when heated in a stream of hydro¬ 
gen. It may, however, be easily prepared, either by the electrolysis 
of fused sodium-aluminium chloride : 2NaCl, Al.^Cl^, or by decom¬ 
posing the chloride with sodium. The latter process can be carried 
out on a small scale by leading the vapour of aluminium chloride 
over fused sodium with a stream of hydrogen. After the violent 
reaction is over, the reduced aluminium is found under the fused 
sodium-aluminium chloride produced at the same time, if the 
sodium was not in excess. 

On the large scale, aluminium is manufactured in the following 
manner. The double chloride of sodium and aluminium is first 
prepared by heating a mixture of alumina, common salt, and char¬ 
coal in a stream of chlorine. 100 parts of the sublimed double 
chloride are next mixed with 35 parts of sodium and 40 parts of 
cryolite (to serve as a flux), and heated on the hearth of a reverbe¬ 
ratory furnace. The aluminium is then reduced, and collects on 
the hearth under the fused slag. 

The high tenacity of aluminium, its property of not rusting, like 
iron, in the air, and its low specific gravity, would make it an ex¬ 
tremely useful metal could it be prepaied in larger quantities and 


Aluminium. 


421 


at a lower price than is at present the case. It is now little used, 
except for scientific purposes and for the manufacture of aluminium 
bronze. 

We are only acquainted with one oxide of aluminium, and only 
with one compound with chlorine and the other halogens. 

Aluminium Oxide, Alumina : A 1 2 0 3 .—This compound occurs 
in nature in crystals belonging to the hexagonal system—partly 
transparent, of a blue or red colour (probably due to traces of cobalt 
and chromium) as the two precious stones, the sapphire and ruby 
*—and partly less pure, as grey, usually opaque, crystals, called 
corundum. Crystalline alumina is distinguished by its hardness— 
next to the diamond it is the hardest substance known, and is, 
therefore, highly valued for grinding and polishing. For this pur¬ 
pose massive crystalline alumina, called emery, is used. Emery 
is found in some parts of Asia Minor, especially on the island of 
Naros, as well as in Massachusetts. 

The oxide may be easily prepared as a white amorphous powder 
by glowing either its hydrate or its double sulphate with ammonium 
—called ammonia alum. 

Like all sesquioxides, alumina is a weak base, and is not soluble 
in acids either in the natural or artificial form. It can only be 
brought into solution by fusion with acid potassium sulphate, or 
with alkaline hydrates or carbonates. Alumina forms salts not 
only with strong acids but also with strong bases, in which it plays 
itself the part of an acid. Some of these salts— e.g. those of potas¬ 
sium and sodium—are soluble in water, hence it follows that 
aluminium hydrate when separated from aluminium salts by caustic 
soda (or potash) dissolves in an excess of the reagent. 

Alumina is infusible at the highest temperatures of our furnaces, 
but melts in the oxy-hydrogen flame to a colourless glass, which 
crystallizes on cooling. By the previous addition of minute quan¬ 
tities of potassium dichromate, artificial rubies may be thus 
obtained. 

Aluminium Hydrate : Al 2 (OH) 6 , or Al(OH) 3 .—This compound, 
occurring in nature as hydrargillite or gibbsite , is obtained as a 
white gelatinous precipitate when a soluble aluminium salt is pre¬ 
cipitated with ammonia. The precipitate, which is difficult to 
wash, dries to a translucent mass, and afterwards forms a light, 
white amorphous powder, insoluble in water. When gently heated 


422 Text-Book of Inorganic Chemistry. 

it loses two molecules of water and is changed into a hydrate of the 
composition : |^Jq which is also found in the mineral king¬ 
dom as diaspore. 

Aluminium hydrate is easily dissolved both by acids and alkalies 
(except ammonia). In both cases salts are produced, the alumina 
playing the part of base or of an acid. It possesses the property 
of uniting with organic colours and precipitating them from their 
solutions, and is, therefore, a valuable mordant. Calico and other 
stuffs, which do not themselves unite with the colours with which 
they are to be dyed, retain them firmly if impregnated with alu¬ 
minium acetate—a salt which readily deposits aluminium hydrate. 
Organic colours precipitated with aluminium hydrate are termed 
lakes. Alumina and its compounds are thus largely used in dyeing. 


Aluminium Salts.— Aluminium, as we have just mentioned, 
forms two classes of salts : those in which it plays the part of a 
base, and those in which it is an acid. The former are the 
common salts of aluminium ; the latter are called aluminates. 
They are found in the crystalline form in nature, and can also be 
prepared artificially. 

Two classes of aluminates are known, one in which the alu??iinic 
acid is tribasic, and one in which it is dibasic. These two acids are 

simply the two hydrates : Al(OH) 3 and j ^Jq The naturally 

occurring aluminates (often called spinelles) belong chiefly to the 
latter class. The following are the more important of these salts. 

Potassium Aluminate : j^lO-OK and sodium aluminate : 

Al(ONa) 3 , are obtained by dissolving aluminium hydrate in caustic 
potash or soda, and separate out on evaporation or on the addition of 

alcohol. Magnesium aluminate —the mineral spinelle : ^Jq j- 0 2 Mg, 


zinc aluminate—gahnite : ^|q j 0 2 Zn, and ferrous ahwiinate — 
iron-spinelle : j 0 2 Fe, are aluminates occurring in nature. 


Aluminium Sulphate: (SO 2 ) 3 O 0 Al 2 + i 8 H 2 0 . — The normal 
salt of the above composition is occasionally found in nature, and 
is prepared by dissolving aluminium hydrate in sulphuric acid, or 


Aluminium Sulphate. 423 

commercially by acting on clay (aluminium silicate), with concen¬ 
trated sulphuric acid. The mixture of aluminium sulphate and 
silicic acid obtained by the latter process is known in trade as 
alum-cake. The salt crystallizes from concentrated solutions in 
fine needles or plates, which deliquesce in the air. Besides the 
normal salt, numerous basic compounds are known. One of these 
is found in nature as the mineral aluminite , and has the composi¬ 
tion : (S 0 2 ) 3 0 e Al 2 + 4 Al(OH) 3 + 21H 2 0 . 

Aluminium sulphate forms crystalline double salts with the 
sulphates of the alkalies, of thallium, and of silver. These double 
sulphates all crystallize in the regular system, and are all isomor- 
phous with one another : they are called alums. The longest 
known of these compounds is potash-alum , potassium-aluminium 

sulphate: |q^} oic' + I2H 2°» or KA 1 (S 0 4 )„ + I2H 2 0, to which 

reference has been previously made (p. 336). It may, like all 
the other alums, be considered as a double sulphate of potas¬ 
sium and aluminium, and its formula would then be : S 0 2 ( 0 K) 2 + 
(SO 2 ) 3 O 0 Al 2 + 2 4 H 2 0 . 

Potash-alum, of the above composition, crystallizes from its 
aqueous solution in large, colourless, and transparent regular 
octahedra, often combined with faces of the cube. If its warm 
solution is mixed with potassium carbonate (or hydrate) so that the 
precipitate at first produced is redissolved, the solution, which 
now contains a basic salt, deposits crystals of ordinary alum in the 
form of cubes, not octahedra. This is called cubic ahem. 

Alum possesses a sweet, astringent taste, has an acid reaction, 
and is easily soluble in water, especially warm. It is insoluble in 
alcohol. When the crystals are exposed to the air they become 
covered with a white opaque crust, which is said to be due to the 
formation of a basic salt from the ammonia of the air, and not to 
loss of water. If alum is heated it melts in its water of crystalliza¬ 
tion, and as the liquid loses water it becomes thicker and thicker 
and froths considerably. Finally, the anhydrous salt remains 
behind as a spongy mass, known in medicine as burnt alum. 

Alum may be prepared directly from its constituents by mixing 
solutions of aluminium and potassium sulphates, and evaporating 
down. The aluminium sulphate used for the manufacture of alum 
is principally obtained in two different ways—either by treating 
kaolin or some other clay (aluminium silicate) with concentrated 


424 Text-Book of Inorganic Chemistry. 

sulphuric acid, or by roasting alum-shale and extracting with water. 
The alum-shale does not contain aluminium sulphate ready formed, 
but consists of a shale (essentially solidified clay) mixed with small 
crystals of iron pyrites, and often with some combustible sub¬ 
stances. When roasted in the air oxidation takes place, with the 
production of ferric oxide (or ferrous sulphate), sulphurous anhy¬ 
dride and sulphuric acid, the latter acid decomposing the alumi¬ 
nium silicate, forming aluminium sulphate. The mass is extracted 
with water, the requisite quantity of potassium sulphate added, and 
evaporated down to crystallize. Some varieties of alum-shale are 
so easily decomposed, and contain the iron pyrites so finely 
divided, that it suffices to moisten them with water and leave them 
exposed to the air for some time to effect the required decomposi¬ 
tion. 

The value of an alum for dyeing purposes is greater the less 
the amount of combined iron it contains, because this substance 
spoils the colour. Alum particularly free from iron is prepared 
from a naturally occurring basic aluminium sulphate, called alunite. 
This substance is much less widely distributed in nature than 
alum-shale, and is chiefly found near Rome, in Tuscany, and 
Hungary. The alum prepared from it is known in trade as 
Roman alum. Alunite, which is insoluble in water, may be con¬ 
sidered as a compound of one molecule of anhydrous alum united 
with two of aluminium hydrate, and is therefore represented by the 

formula: |o 2 joK^ + 2A1 ( 0H )s- When gently heated, it yields 

water, insoluble alumina, and soluble potash alum. The last 
named compound is then extracted with water, and usually crystal¬ 
lizes in cubes. 

If a solution of aluminium sulphate is mixed with one of 
sodium, ammonium, rubidium, caesium or thallium sulphate, alums 
are produced of the same crystalline form as potash-alum and of 
analogous composition. An alum of silver has also been obtained. 
These different alums are distinguished by the names soda-, 
ammonia-, &c., alum. Formerly, large quantities of ammonia- 
alum were manufactured in England from the ammoniacal liquors 
of the gas-works ; the manufacture has now almost entirely ceased, 
owing partly to the immense quantities of cheap potassium 
compounds obtained from the Stassfurt beds. Ammonia-alum 
cannot be distinguished from potash-alum in its external cha¬ 
racters. 


Alum . 


425 


Besides the substitution of the potash by other monoxides, the 
alumina in alum may also be displaced by other sesquioxides with¬ 
out altering the shape or general character of the compounds ; and 
alums can be prepared in which the aluminium is displaced by iron, 
manganese, or chromium respectively. These alums are known by 
the names of the triad metals which they contain, and have the 
following composition :— 



Iron-alum 


Manganese-alum 


Chrome-alum . 


These alums are coloured : iron-alum usually violet, chrome- 
alum, dark-green or purple. Like alum proper, the potassium 
can be displaced by other monad metals, and we are thus ac¬ 
quainted with more than twenty isomorphous compounds belong¬ 
ing to this class. 

Aluminium Phosphate : PO - 0 3 A 1 + 4H 2 0.—The normal salt 
is deposited as a white gelatinous precipitate, when a solution of 
sodium phosphate is added to one of aluminium sulphate or of 
alum. It is easily soluble both in acids and alkalies (except 
ammonia). Various basic phosphates are known, among which is 
the mineral wavellite . 

Aluminium Silicate : (SiO) 3 0 6 Al 2 .—A silicate having this 
composition is thrown down as bulky precipitate on adding excess 
of a hot dilute solution of aluminium sulphate to a neutral solu¬ 
tion of sodium silicate. A salt of the composition : Al 2 0 3 ,Si 0 3 
is found crystallized in the mineral kingdom as andalusite and 
disthene. In combination with other silicates, aluminium silicate 
is one of the most important constituents of the crust of the earth. 
Among these may be mentioned the felspars of which the com¬ 
position has been already given (p. 272). The commonest of these 
is potash-felspar ox orthoclase\ KAlSi 3 0 8 , less common axe soda- 
felspar or albite , of similar composition, and lime-felspar or anor- 
thite. Most of the naturally occurring double silicates are anhy¬ 
drous, but some—the zeolites — contain water. To this class 
belong the minerals natrolite and analcime , which are hydrated 


426 Text-Book of Inorganic Chemistry. 

double silicates of aluminium and sodium. They belong to the 
few silicates which are decomposed by hydrochloric acid. The 
felspars and most of the other silicates are quite unacted upon 
by hydrochloric acid, and can only be decomposed by treatment 
with hydrofluoric acid, or by fusion with sodium (or potassium) 
carbonate. 

One of the most important hydrated silicates is clay , which in 
its purest form is kaolin or china-clay. This is hydrated alu¬ 
minium silicate of the approximate composition : A1 2 0 3 , 2Si0 2 + 
2H 2 0. It is produced by the gradual weathering of felspar, which 
is often found mixed with it. The process consists in a gradual 
change of the potassium silicate by the carbonic acid of the air into 
soluble potassium carbonate (see p. 344), leaving insoluble alu¬ 
minium silicate, which then unites chemically with water. 

Kaolin is largely used for the manufacture of porcelain and 
earthenware. Clay when finely powdered and mixed with water 
forms a plastic mass which readily allows itself to be moulded into 
various shapes, and if the clay is afterwards heated, it bakes 
together to a hard but porous mass. It is because of these two 
properties that clay is so largely used for the manufacture of articles 
of china, stoneware, and earthenware. But vessels prepared from 
clay alone are porous, and cannot be used to contain liquids : to 
make them available for this purpose they are either coated with 
a thin layer of glass (the glaze ), or else the clay is previously mixed 
with some fusible material (the frii), then burnt and glazed. In 
the manufacture of true porcelain the frit employed is usually 
either felspar or a mixture of felspar and quartz, which is carefully 
mixed with the kaolin and then burnt at a high temperature. In 
true porcelain, the frit penetrates through the entire mass—giving 
it its well-known translucent appearance. Porcelain is usually 
glazed by dipping it in finely powdered felspar suspended in water, 
and then again heating. 

Stoneware of very various qualities is a coarser kind of porce¬ 
lain made with a more easily fusible frit which does not penetrate 
through the entire mass of the clay. Stoneware is glazed with a 
more fusible glaze than felspar. 

Earthenware or Fayence , which is made from variously 
coloured clays, is porous, and of a yellow or reddish colour. It is 
glazed either with lead oxide, forming an easily fusible lead silicate, 
or with common salt. The salt is usually thrown into the kiln 


Ultramarine. 


427 


when the burning is nearly over. It then volatilizes, and is de¬ 
composed with the water-vapour present and the silica of the clay, 
forming hydrochloric acid and a sodium silicate, which acts as a 
glaze. 

A clay rich in calcium carbonate, and which is sometimes mixed 
with clay in making earthenware, is known as rnarl. Intimate 
mixtures of clay with ferric oxide of various colours from yellow to 
brown are used as pigments under the name of ochres. Fullers 
earth and bole are fine clays used for cleansing clothes and other 
purposes. 

Ultramarine.— The highly valued ornamental stone called 
lapis lazuli , and occurring especially in China and Thibet, was 
previously powdered and sold under the name of ultt'amaruie for 
its weight in gold. This substance is essentially a compound of 
sodium and aluminium silicates with sodium pentasulphide. Fifty 
years ago Gmelin discovered a method for the artificial manufac¬ 
ture of ultramarine, and since then his processes have been so 
perfected, that the artificial substance can now be made of a finer 
colour than that found naturally. This discovery has, therefore, 
been well described as one of the most brilliant victories of chemical 
science during the present century. 

The artificial preparation of ultramarine appears at first ex¬ 
tremely simple, but in order to prepare a pigment of fine colour a 
number of practical details must be attended to. These can only 
be acquired by long experience, and are often preserved as trade 
secrets. 

Ultramarine is obtained by heating clay in a closed crucible 
with some mixture which produces sodium sulphide, such as 
sulphur and sodium carbonate, or sodium sulphate and charcoal. 
The product first produced is green—ultramarine-green—but if 
mixed with sulphur and again gently heated, changes to the well- 
known blue colour of ultramarine proper. The change in colour 
is probably due to the formation of a polysulphide of sodium. If 
heated too strongly, the colour is destroyed. 

The chemical constitution of ultramarine is still uncertain. It 
is thought that the blue compound contains sodium pentasulphide, 
and the green the disulphide. The sodium sulphide cannot be dis¬ 
placed by potassium sulphide. The products are finely ground, 
brought into trade in various tints, and largely used for many 
purposes. 


428 Text-Book of Inorganic Chemistry 

Ultramarine is not so permanent as smalt (cobalt blue), nor as 
indigo, but is not decomposed by soaps and alkalies like Prussian 
blue. Acids, however, at once decompose ultramarine, and it may 
be recognized by the reaction that when mixed with hydrochloric 
acid, sulphuretted hydrogen is evolved, much white sulphur sepa¬ 
rates—from the sodium pentasulphide—and the blue colour en¬ 
tirely disappears. 

Aluminium Chloride : A1 2 C1 6 .—This compound may be ob¬ 
tained by burning the metal in a stream of chlorine. It is, 
however, usually prepared by intimately mixing alumina and 
powdered charcoal (best sugar charcoal*) with starch paste, and 
then proceeding as described under the preparation of silicon 
tetrachloride (see p. 273 ). On glowing in a stream of dry chlorine, 
aluminium chloride sublimes in yellowish crystalline plates. The 
yellow colour is due to traces of ferric chloride, which may be 
removed by sublimation over hot aluminium, when the volatile 
ferric chloride is converted into non-volatile ferrous chloride. 
Pure crystalline aluminium chloride is colourless, sublimes when 
gently heated, and rapidly attracts water from the air. 

The aqueous solution of the chloride, which may be obtained 
by dissolving aluminium hydrate in hydrochloric acid, when 
evaporated down deposits colourless deliquescent crystals of the 
composition: A1 2 C1 6 + i 2 H a O. The water cannot be expelled 
from these crystals by heating so as to leave the anhydrous 
chloride, because the aluminium chloride is then decomposed with 
the formation of hydrochloric acid, and a basic chloride or even 
alumina. Aluminium chloride unites with sodium and other 
chlorides, producing fairly stable double salts. 

Aluminium Fluoride : A1 2 F 6 , is obtained by dissolving alumi¬ 
nium hydrate in hydrofluoric acid, or by glowing a mixture of 
alumina and fluor-spar in hydrochloric acid gas. It is insoluble in 
water, sublimes at a high temperature in colourless cubes, and in 
combination with sodium fluoride forms the mineral cryolite : 
Na 3 AlF 6 = 6 NaF,Al 2 F 6 . 

Aluminium Sulphide'. A1 2 S 3 , is produced when alumina is • 
glowed in the vapour of carbon disulphide. It is a bright yellow 
vitreous mass, which with water, or even moist air, rapidly decom¬ 
poses into sulphuretted hydrogen and aluminium hydrate. 

1 Or a very pure form of lamp-black, known as gas-black. — Ed. 


Beryllium. 


429 


Detection of Aluminium Compounds. 

Aluminium salts are distinguished from those of the alkalies 
and alkaline earths by the production of a white gelatinous pre¬ 
cipitate of aluminium hydrate with ammonia—a precipitate which 
is insoluble in ammonium salts and is therefore produced in their 
presence (comp, magnesium). They are further distinguished »by 
yielding a white precipitate of aluminium hydrate, not sulphide, 
with ammonium sulphide. 

Like all sesquioxides, alumina is a weak base, and does not 
form compounds with weak acids— e.g. carbonic acid. If a solu¬ 
tion of aluminium chloride is mixed with one of sodium carbonate, 
aluminium hydrate and sodium chloride are formed, and carbonic 
acid set free. If the sodium carbonate is in excess, acid sodium 
carbonate is produced. The weak basic properties of alumina are 
further shown by its behaviour with caustic soda (or potash). 
Caustic soda precipitates aluminium hydrate from solutions of 
aluminium salts, but the precipitate dissolves in an excess of the 
alkali, forming sodium aluminate—a salt in which alumina plays 
the part of an acid (p. 423 ). Aluminium compounds when heated 
on charcoal before the blowpipe leave a white infusible mass, 
which, when moistened with a drop of cobalt nitrate and again 
heated becomes of a fine blue colour. 


BERYLLIUM. 

Chemical Symbol : Be .—Atomic Weight'. 13 - 6 . 

This somewhat rare element is chiefly found in nature as a 
double silicate of aluminium and beryllium : A1 2 0 3 , Be 2 0 3 ,6Si0 2 , or 

(SiO ) 3 which, in its rough opaque form, is called beryl; 

when purer and transparent it is either of a pale sea-green colour : 
aquamarine , or else of a bright green : emerald , both of which 
varieties have considerable value as precious stones. It further 
occurs as silicate simply in the mineral phenakite : 2 Be 2 0 3 , 3 Si 0 2 , 
or Si 3 (0 3 Be) 4 , and as double oxide of aluminium in chrysoberyl: 
Be 2 0 3 , A1 2 0 3 . 

The metal may be obtained by leading the vapour of the 



430 Text-Book of Inorganic Chemistry . 

chloride over fused sodium contained in iron boats, and afterwards 
washing with water. It is a white metal like aluminium, with a 
specific gravity of 17 , and melts at a red-heat under partial oxida¬ 
tion. It is without action on water at the ordinary temperature, 
but decomposes it slowly on boiling. Dilute hydrochloric or 
sulphuric acid easily dissolves it with rapid evolution of hydrogen. 
Concentrated nitric acid scarcely acts on it. It dissolves also in 
the caustic alkalies (not in ammonia), forming soluble compounds 
resembling the aluminates. 

Beryllium Oxide or Beryllia : Be 2 0 3 , is a white, amorphous 
powder of 3-0 specific gravity, obtained by heating 

Beryllium Hydrate'. Be 2 (OH) 6 .—This compound is thrown 
down as a white gelatinous precipitate when ammonia is added to 
a solution of a beryllium salt. The precipitate forms a white 
powder when washed and dried; it is soluble in caustic potash 
or soda and in ammonium carbonate, from which solution it is 
again precipitated on boiling as a carbonate. 

The salts of beryllium possess a pure sweet taste, whence the 
name glucinum , which is sometimes given to the metal; they are 
mostly soluble in water. 

Beryllium Sulphate : (S0 2 ) 3 0 6 Be 2 + I 2 H 2 0 , easily crystallizes 
when its aqueous solution is evaporated down. It forms a double 
compound with potassium sulphate, which is difficult to obtain in 
the crystalline form, and has a different composition to common 
alum, with which it is therefore not isomorphous. 

Beryllium Chloride : Be 2 Cl 6 , is prepared like aluminium chloride, 
which it resembles, but is less volatile. It forms, when pure, white, 
crystalline, highly deliquescent needles. 

Beryllium Carbonate. —On adding sodium carbonate to a solu¬ 
tion of a beryllium salt a precipitate of a basic carbonate is pro¬ 
duced, which dissolves with difficulty in an excess of the reagent. 
This carbonate is much more easily soluble in ammonium carbo¬ 
nate, a fact which enables us to detect beryllium and to separate 
it from aluminium. Water containing carbonic acid also slightly 
dissolves the basic carbonate. 

From the above properties of beryllium and its compounds it 
will be seen that the metal occupies a peculiar position between 
magnesium and aluminium, both of which it resembles in some 


Gallium, 


431 

points. And it has long been a matter of discussion whether the 
element should be classed with the alkaline-earths or with the 
earths. In the former case it would be a dyad with an atomic 
weight of 9*1, and its oxide would be written : BeO ; in the latter 
the element is a triad, its atomic weight is half as much again, or 
13*6, and its oxide becomes : Be 2 O a . Recently the specific heat of 
the metal has been determined and found to be 0*445, which when 
multiplied by the atomic weight (13*6) gives 6*05, proving, according 
to Dulong and Petit’s law (p. 70), that this is the true atomic weight 
and not 9*1. 


GALLIUM. 

Chemical Symbol : Ga.— Atomic Weight : 69*8. 

This element is contained in some samples of blende found in 
the Pyrenees, but only in extremely minute quantities (about 0*002 
per cent.) Its spectrum contains two characteristic lines in the 
violet which led to its discovery by the French chemist, Lecoq de 
Boisbaudran in 1875. 

Gallium is a white lustrous metal of low melting-point (30°); 
when once melted it remains liquid, like mercury, even at o°. Its 
specific gravity is 5*9. The metal remains unaltered in the air, 
and only becomes covered with a thin layer of oxide when heated 
nearly to redness. It is not attacked by water at the ordinary 
temperature, but dilute hydrochloric acid as well as the alkalies 
dissolve it with evolution of hydrogen. 

The compounds of gallium resemble those of aluminium. Gal- 
lium oxide : Ga 2 0 3 and the hydrate are white amorphous powders, 
insoluble in water. Gallium chloride : Ga 2 Cl 6 , which may be ob¬ 
tained by heating the metal in a stream of dry chlorine, is soluble 
in water and easily sublimes. A dichloride : GaCl 2 is also known. 

Ammonium-gallium alum : gQ 2 J ONH + crystallizes 

in regular octahedra. 




432 


Text-Book of Inorganic Chemistry 


INDIUM. 

Chemical Symbol : In.— Atomic Weight'. 113 ' 4 - 

Indium, like gallium, belongs to the rarest metals ; it is also 
found in very small quantities in some zinc ores, and when these 
are worked for zinc alloys itself with this metal. The zinc from 
Freiberg in Saxony contains from 0-05 to o-i per cent, of indium. 
If such zinc is treated with insufficient hydrochloric acid to com¬ 
pletely dissolve it, the indium remains behind with other metals, 
and is purified by different processes. 

Indium is a white lustrous metal, of specific gravity 7-4, and 
melting at 176°. It dissolves slowly in hydrochloric or dilute sul¬ 
phuric acid, but easily in nitric acid. The metal remains unchanged 
in the air, but burns, when heated to redness, with a violet light, 
and producing brown vapours. Indium is a triad in all its com¬ 
pounds. 

Indium Oxide : ln 2 0 3 , is a bright yellow powder, which becomes 
dark brown when heated. 

Indium Hydrate'. In(OH) 3 , is thrown down from indium salts 
on addition of ammonia, as a white gelatinous precipitate, resem¬ 
bling aluminium hydrate; it also dissolves in caustic soda or 
potash. 

Indium Sulphate : (S 0 2 ) 3 0 6 ln 2 , remains as a gummy mass 
when its solution is evaporated down. It unites with ammonium 
sulphate, but not with potassium sulphate, and forms ammonium- 

indium alum: sq 2 }oNH + i2 HoO. 

Indium Chloride : InCl 3 , is prepared by heating either the metal 
or a mixture of the oxide and charcoal in a stream of dry chlorine, 
when it sublimes in white crystalline plates. It is very deliques¬ 
cent. 

Indium Sulphide : In 2 S 3 , is produced when sodium and sulphur 
are gently heated together, and then forms a brown infusible mass. 
Sulphuretted hydrogen precipitates the same compound from neu¬ 
tral solutions of indium salts, or when sodium acetate has been 
added to the solution. In the presence of free hydrochloric acid 
no precipitate is obtained. Indium, which is otherwise closely allied 
to aluminium and gallium, in this respect resembles zinc. 

Traces of an indium compound, if introduced in the non- 
luminous gas flame, may be at once detected by its spectrum, 


Yttrium , Terbium , Erbium , and Ytterbium. 433 

which is characterized by one line in the dark blue. It was this 
line which led to the discovery of the metal, and from which it 
derives its name. 


YTTRIUM, TERBIUM, ERBIUM, and YTTERBIUM. 

Chemical Symbols : Y, Tr, Er, Yb. 

Atomic Weights : 90, 148, 166, 173. 

These four rare elements, which have not yet been prepared in 
the free state, are only found in a few extremely scarce minerals— 
e.g. in gadolinite as silicates, in ytirotantalite and samarskite as 
tantalates and niobates. 

Of their oxides, yttrium and ytterbium oxides: Y 2 0 3 and 
Yb 2 0 3 , are white, terbium oxide : Tr 2 0 3 , is yellow, and erbium 
oxide : Er 2 O g , pink. 

The salts of all four so closely resemble one another that they 
can only be partially separated with considerable difficulty. The 
sulphates and nitrates are soluble in water, the former compounds 
being more soluble in cold than hot water. The chlorides are non¬ 
volatile, and can be prepared anhydrous like magnesium chloride. 
Finally, the oxalates and carbotiates are insoluble, so that they can 
be separated from aluminium and iron (in the ferric state) by pre¬ 
cipitating with oxalic acid. 

The salts of erbium are of a pink colour, and yield a very charac¬ 
teristic absorption spectrum, by means of which small quantities of 
the element may be detected. 


CERIUM, LANTHANUM, and DIDYMIUM. 

Chemical Symbols : Ce, La, Di. 

Atomic Weights : 141, 139, 146. 

These three elements, which may be prepared by the electro¬ 
lysis of their fused chlorides, always occur associated with one 
another in nature—usually as a hydrated silicate : cerite , which is 
decomposed by hydrochloric or sulphuric acid. 

Cerium forms two oxides —viz. a sesquioxide : Ce 2 0 3 , and a per¬ 
oxide : Ce 0 2 . Both are white compounds, but the latter becomes 




434 Text-Book of Inorganic Chemistry. 

yellow when heated in the air. The two other oxides, lanthanum 
oxide : La 2 0 3 , and didymium oxide : Di 2 O a , correspond to the ses- 
quioxide of cerium. Didymium also forms a brown peroxide. 

The salts of the three metals resemble one another so closely that 
it is difficult to separate them. The cerium may be got rid of as the 
peroxide when the alkaline solution containing cerium hydrate in 
suspension is acted on by chlorine, but the separation of lantha¬ 
num and didymium is much more difficult. 

The sulphates , nitrates , and chlorides closely resemble those of 
the preceding group. The oxalates are also insoluble in water 
and dilute acids. The sulphates are distinguished by forming 
a double sulphate with potassium (not an alum), which is quite 
insoluble in a saturated solution of potassium sulphate, and so 
enables us to separate this group of metals from the preceding 
group. 

Didymium compounds are pink, like those of erbium, but with 
a bluer tinge. They also yield a characteristic absorption spec¬ 
trum, of which the darkest band is one in the yellow. 


THORIUM. 

Chemical Symbol : Th.— Atomic Weight’. 232’ 

This extremely rare element is only found in two scarce minerals, 
both of which are silicates—the one anhydrous, called thorite , and 
the other hydrated, orangite. 

The metal can be extracted by acting on its volatile chloride : 
ThCl 4 , with sodium. The only oxide has the composition : Th 0 2 . 
Its sulphate , chloride , and nitrate are all soluble in water. 


ZIRCONIUM. 

Chemical Symbol : Zr.— Atomic Weight : 90. 

Zirconium is also one of the rarer elements, and chiefly found 
in its silicate, the minerals zircon and hyacinth : Si( 0 4 Zr). The 
metal is obtained in brilliant brittle scales, by strongly heating 
potassium fluozirconate with aluminium in a graphite crucible, and 




Zirconium. 


435 


afterwards extracting the excess of aluminium with hydrochloric 
acid. It has a specific gravity of 4*1, is very difficultly fusible, and. 
is not attacked by oxygen at a red heat, but burns in the oxy-hvdro- 
gen flame. Hydrochloric, nitric, or sulphuric acid scarcely attacks 
it, but it dissolves in hydrofluoric acid with evolution of hydrogen. 

The mineral zircon is not attacked by any acid, but is decom¬ 
posed by fusing it either with sodium carbonate, or, better, with 
acid potassium fluoride ; in the latter case potassium fluosilicate 
and fluozirconate are formed, of which the latter dissolves when the 
fused mass is boiled with water, while the former remains undis¬ 
solved. From this solution ammonia precipitates 

Zirconium Hydrate : Zr(OH) 4 , as a white voluminous mass, 
which loses water when heated, and leaves zirconium oxide or 
zirconia : Zr 0 2 , as a white power, insoluble in dilute acids. 
Caustic soda also precipitates the hydrate like ammonia, which is 
soluble in an excess of this reagent like aluminium and beryllium 
hydrates. 

Zirconium Sulphate'. gQ 2 *Q 3 |zr, obtained by dissolving the 

hydrate in sulphuric acid, crystallizes with difficulty from its aqueous 
solution. 

Zirconium Chloride : ZrCl 4 , prepared like aluminium chloride, 
by glowing a mixture of zirconia and charcoal in dry chlorine, is a 
white crystalline substance, volatile without decomposition, and 
soluble in water. The same solution is obtained by dissolving 
zirconium hydrate in hydrochloric acid, from which, however, an 
oxychloride : ZrOCl 2 , with varying quantities of water, crystallizes 
in colourless silky prisms on concentration. Zirconium fluoride : 
ZrF 4 , sublimes as a colourless crystalline mass when an intimate 
mixture of powdered zircon and fluor-spar is heated in a stream of 
hydrochloric acid. It is insoluble in water and hydrochloric acid, 
but dissolves in hydrofluoric acid. On mixing this solution with 
potassium fluoride and evaporating down, crystals of potassium 
fluozirconate : ZrF 4 + 2KF, or K 2 ZrF 6 are obtained, isomorphous 
with the corresponding compounds of silicon, titanium, and tin. 

Zirconium (like thorium) is a tetrad in nearly all its com¬ 
pounds ; it is connected with silicon and titanium on the one hand, 
and with tin on the other, as is readily seen from the above 
description of its compounds. 


f f 2 


43^ 


Text-Book of Inorganic Chemistry. 


HEAVY METALS. 

IRON. 

Chemical Symbol : Fe .—Atomic Weight'. 56. 

Iron belongs to the most widely distributed, and to the most 
important elements. It is found in the free state in meteorites 

_larger or smaller masses of iron which occasionally reach the 

earth from extra-terrestrial space. These meteorites always contain 
nickel, often as much as 10 per cent., and may be distinguished 
from metallic iron of other origin by peculiar markings (first noticed 
by Widmanstadt) which are produced when the surface is polished 
and treated with nitric acid. 

The most important" compounds of iron, and those used ex¬ 
clusively for the extraction of the metal, are its oxygen compounds. 
They are not only very widely distributed, but are often found in 
immense quantities. Among these are : specular-iron, ferric oxide : 
Fe 2 0 3 , beautifully crystallized in the hexagonal system, and espe¬ 
cially abundant on the Island of Elba. Ferric oxide, as granular, 
columnar, or earthy masses forms the important ore, red hcematitc, 
which when combined with water is known as br'own hcematite or 
limonite. Red haematite is principally found in Lancashire and 
Cumberland in England, while brown haematite is found either » 
associated with carboniferous rocks in the Forest of Dean and 
Glamorganshire, or else with oolitic rocks, as the earthy haematite 
of Northamptonshire and Lincolnshire. The bog iron-ores of 
Ireland and North Germany also consist of limonite, but of a 
much more recent date. Magnetic iron ore : Fe 3 0 4 , when pure, is 
the richest, and one of the most valuable of iron ores. It is found 
crystalline, but usually in rough masses, or as sand, and is the 
chief ore from which the Swedish iron is extracted. Spathic iron 
ore or siderite is impure ferrous carbonate: CO- 0 2 Fe, and is 
found in small quantities as such in England, but chiefly occurs in 
some German localities, and in Styria. Spathic iron ore is gene¬ 
rally found in England, intimately associated with clay as clay iron¬ 
stone, or argillaceous iron ore —one of the most important, if not 


Iron. 


437 


the most important English ores of iron. It occurs in large quan¬ 
tities in various parts of the kingdom. 

Other compounds of iron which are found in large quantities in 
nature but which cannot be used as iron ores are : Iron pyrites : 
FeS 2 , distinguished by its bright yellow colour and the readiness 



Fig. 65. 

with which it crystallizes; magnetic pyrites : Fe 3 S 4 or Fe 7 S 8 , richer 
in iron; copper pyrites : CuFeS 2 ; arsenicalpyrites : FeAsS, &c. 

Iron is extracted from its ores by reduction with carbon at a 
high temperature. The furnaces used for iron smelting, called 
blast-furnaces , are usually built in the form represented in fig. 65. 
They are of strong masonry, lined with fire-bricks, and reach a 
height of sixty or eighty feet. The iron ore, which, if clay iron ore, 

























438 Text-Book of Inorganic Chemistry. 

must first be roasted to expel carbonic acid and water, is introduced 
at the top with alternate layers of coke, or in some countries 
(Sweden) charcoal. And if the ore does not contain the requisite 
constituents to form a fusible slag, either limestone or silica in 
some form must be added for this purpose. At the lower part of 
the furnace a blast of air is blown in from several iron openings, 
called tuyeres , which are kept at a low temperature by a stream of 
cold water circulating through them. Above the blast about D in 
the figure, the iron is reduced to a spongy mass, partly by the hot 
carbonic oxide, and partly by the free carbon. This reduced iron 
takes up carbon, &c., in its descent, becomes more fusible, and 
finally forms a liquid mass on the hearth at G. At the same time 
the silica, lime, alumina, &c., unite to form a fusible slag (p. 404), 
which is lighter than the iron, and by floating on its surface pro¬ 
tects it from the oxidizing action of the blast. When the furnace 
is at work, a continuous stream of molten slag flows out by a side 
opening—over a raised portion—the dam-plate (not shown in the 
figure). The inflammable gases-which escape from the top of 
the furnace consist chiefly of carbonic oxide, produced by the 
reduction of the carbonic acid, and of nitrogen. In modern 
furnaces these gases are not allowed to burn in this wasteful 
manner at the mouth of the furnace, but pass out at a side opening 
near the top, and are then used to heat the blast of air entering at 
the base of the furnace. A large economy of fuel is effected by 
using a hot instead of a cold blast. In such furnaces, the top is 
closed with a valve arrangement (cup and cone ) through which the 
materials are introduced when necessary. At stated intervals, 
when sufficient molten iron has collected on the hearth G, a hole 
at the base, which is closed with sand or clay, is opened with 
a long iron rod, and the iron flows out into suitable moulds, pro¬ 
ducing rough bars—called pigs. Such a furnace is kept in con¬ 
tinuous action for several years. 

Cast-iron obtained in this way is still very impure ; it may be 
purified from the slag which it contains mechanically by re-melting 
with a suitable flux. Even then it contains small quantities of 
sulphur, phosphorus, and of silicon reduced from the slag, as well 
as up to four per cent, of carbon, partly as graphite and partly in 
chemical combination with the iron. Two varieties of cast-iron 
are distinguished—viz. white cast-iron and grey cast-iron. The 
former is generally produced when the temperature of the furnace 
is low, and when the metal is rapidly cooled. It contains nearly 


Iron. 


439 


the whole of its carbon combined with the iron, is silver white 
in colour, whiter than pure iron, and with a lamellar, crystalline 
fracture. The latter variety—grey cast-iron—is produced with 
a high temperature and slow cooling, and contains most of its 
carbon in the state of graphite. Grey cast-iron is darker in colour 
than white iron, and has a granular crystalline fracture. Mottled 
cast-iron is an intermediate variety between these two forms. 
When cast-iron is dissolved in dilute hydrochloric acid, the carbon 
present as graphite remains behind, while that in chemical com¬ 
bination with the iron combines with the nascent hydrogen and 
forms volatile hydrocarbons. Spiegeleisen and ferro-manganese are 
varieties of cast-iron containing manganese, and are largely used 
for the manufacture of steel. 

Cast-iron melts at a lower temperature than any other form of 
the metal, but is brittle, and cannot be worked up into any shape. 
It is, therefore, only used for those articles which can be cast in 
moulds. But as the percentage of carbon diminishes the malle¬ 
ability increases, and the iron becomes more difficultly fusible. 
Iron containing the smallest quantity of carbon (not more than 
0-5 per cent.) is called wrought-iron , while intermediate between 
-wrought- and cast-iron are the various kinds of steel. In order, 
therefore, to convert cast-iron into wrought-iron it is necessary to 
abstract the greater part of the carbon which it contains. This is 
done by two processes, both depending on the same principle, 
called puddling and fining. In the former process, which is most 
generally used in England, the crude pig-iron is melted with slags 
rich in ferric oxide on the hearth of a reverberatory furnace with side 
openings. A considerable quantity of its silicon and carbon is re¬ 
moved by contact with the hot air blowing over it, and by the oxy¬ 
gen of the slag. Long iron bars are then inserted by the workmen 
at the side doors, and the molten metal stirred and worked up so 
as to expose it still further to the oxidizing action of the air. As 
the process goes on, the iron gradually becomes more difficultly 
fusible, and at last forms a tough mass, in which state it is removed 
from the furnace and hammered to remove the fusible slag which 
it still contains. It is then further hammered or rolled into bars 
or plates. 

In the process of fining , the cast-iron is first melted on a hearth,' 
and then a blast of air blown through it from several tuyeres. A 
portion of the iron is oxidized to the magnetic oxide, and nearly 
the whole of the carbon and silicon is removed. After the mass 


440 Text-Book of Inorganic Chemistry. 

has become pasty it is worked, like puddled iron, by hammering 
apd rolling. 

Wrought-iron prepared in either of these ways is never quite 
free from carbon, but only contains from oi to 0*5 per cent. Traces 
of silicon are also usually present, and, what is of more importance, 
traces of sulphur and phosphorus. Iron containing even a small 
quantity of sulphur is brittle when hot, or is said to be red-short , 
while that containing phosphorus to the extent of more than a half 
per cent, is brittle when cold, or, as it is said, is cold-short. 

Wrought-iron has a pale grey colour, is very tenacious, ductile, 
and malleable. Its specific gravity is about 7-8. At a red-heat it be¬ 
comes pasty, and can then be readily welded together and fashioned 
into various shapes. The smith, when he wishes to weld together 
two pieces of iron, makes them both red-hot in the forge, sprinkles 
them with sand, and then unites them into one solid piece by 
hammering. Without the addition of sand or some such substance 
(borax), the two surfaces of iron would be unable to unite with one 
another, because of the layer of oxide with which they become 
coated when heated. The silica of the sand combines with the 
iron oxide, forming an easily fusible slag, which is removed by 
hammering, leaving the two surfaces of bare metal, which unite • 
together under the pressure of the blows. 

Wrought-iron, during its process of manufacture, acquires a 
fibrous structure, which greatly increases its resistance to breaking 
or tearing. And it is because of this property that wrought-iron is 
used for so many purposes for which brittle cast-iron, with its 
crystalline, granular structure, is unsuitable. An axle of a wheel, 
for example, if made of cast-iron, would break after a few turns on 
a rough road, but one of wrought-iron stands for years ; and if it 
occasionally happens that an axle of a railway carriage breaks, it 
is always found that the fibrous wrought-iron has gradually become 
granular and crystalline by the long-continued vibrations. 

Steel is a modification of iron intermediate between cast- and 
wrought-iron both in its properties and composition, as far as carbon 
is concerned. Steel contains about 1 -o per cent, of combined carbon; 
it unites the hardness and fusibility of cast-iron with the ductility 
and malleability of wrought-iron. Steel can be produced either 
from wrought- or cast-iron. The old, and up to twenty-five years 
ago the only known, method of obtaining steel was from wrought- 
iron, by imbedding bars of the metal in charcoal-powder contained 
in closed iron boxes, and then heating to bright redness for about 


Iron. 


441 


a week. By this process carbon is gradually taken up by the iron 
in some unknown manner, and the latter becomes completely con¬ 
verted into steel. The process is called ceme?itation , and the product 
cement-steel. Cement-steel can be rendered more uniform in struc¬ 
ture by working under the hammer at a red-heat, or by fusing and 
casting into blocks. It then forms what is called cast-steel. 

Steel can now be much more economically obtained directly 
from cast-iron. Two processes are chiefly employed to effect this 
change—known as the Bessemer and Siemens-Martin processes 
respectively. In the Bessemer process five to ten tons of cast-iron 
are fused and poured into the converter —a large egg-shaped vessel 
lined with fire-clay, with a narrow mouth and with an opening at 
the base through which air can be blown. A current of air is then 
blown through the molten metal, which oxidizes nearly the whole of 
the carbon and silicon. As soon as this point is reached the blast 
is stopped, a small quantity of spiegeleisen is added to impart car¬ 
bon and manganese, and the metal at once run into moulds. The 
cast-iron employed to manufacture Bessemer steel must be free 
from sulphur and phosphorus if a converter with the usual lining 
of siliceous fire-clay is used, as these impurities are not removed 
during the process. It has, however, been recently discovered that 
if the lining of the converter contains free lime, this substance 
unites with the sulphur and phosphorus and a steel free from these 
impurities is obtained, even when an impure cast-iron is used. Steel 
is also made from cast-iron by the Siemens-Martin process, which 
consists in fusing cast-iron in a specially constructed furnace, and 
adding wrought-iron to it until the requisite composition has been 
obtained. Some spiegeleisen is then added as in the Bessemer 
process, and the steel cast into moulds. 

Steel is of a pale grey colour and takes a high polish. It is more 
easily fusible than wrought-iron, and its structure is finely granular, 
not fibrous, in consequence of which it is more brittle than wrought- 
iron. Steel is especially distinguished from wrought-iron by its 
hardness, which it especially acquires when heated to redness and 
rapidly cooled in water; it is then also exceedingly brittle and 
highly elastic. Glass-hard steel—i.e. steel hardened in this way, 
has but few uses ; in order to adapt the steel to the wants of daily 
life it must be toughened or tempered. This is done by again 
heating and again rapidly cooling, and the higher the temperature 
of the second heating up to a certain limit, the tougher and less 
hard the steel. If the steel is polished it gradually becomes 


44 2 Text-Book of Inorganic Chemistry. 

covered with a thin coating of oxide, the colour of which varies 
from a pale yellow to a dark blue, according to the temperature 
(240° to 300°) to which it has been heated. Cutting tools of various 
kinds are tempered in this way according to the purpose for which 
they are to be used. The specific gravity of steel is less after 
hardening (7-6) than it is before (7-8). 

The preparation of chemically pure iron in the compact state is 
exceedingly difficult, if not impossible. If pure iron is fused in a 
crucible of gas-carbon it combines with some of the carbon ; if 
in one of porcelain, it reduces small quantities of silicon and 
aluminium with which it unites chemically ; and approximately 
pure iron can only be obtained by fusing in a lime crucible with 
the oxy-hydrogen blowpipe. Pure iron in the form of a powder 
may be obtained by reducing pure ferric oxide or ferrous chloride 
in a stream of hydrogen. It then remains as a dark grey powder 
(reduced iron), which is so finely divided that it takes fire in the air 
when it has not been strongly heated in the stream of hydrogen. 

Iron in all its forms is strongly attracted by the magnet: we 
say it is magnetic. Its property in this respect far exceeds that of 
any substance, and is only feebly approached by the magnetic 
power of the metals nickel and cobalt. Soft iron becomes mag¬ 
netic when brought near a magnet, but loses its magnetism as soon 
as the magnet is removed. Steel, on the other hand, does not so 
readily become magnetised, but retains its magnetism for some 
time. Artificial magnets are made of straight or curved steel bars. 
A steel magnet gradually loses its magnetism when warmed, but 
partially regains it on cooling. If, however, the magnet is heated 
to redness it permanently loses its magnetism. 

Notwithstanding the strong attraction of iron for oxygen, the 
metal remains unchanged in dry air and oxygen, but moist air, 
and especially water containing dissolved air, gradually oxidizes it 
—we say, the irpn rusts. A drop of water placed on a polished 
strip of iron and allowed to evaporate leaves behind a dull reddish 
stain of oxide— iron-rust. This rust is not the oxide but the hydrate, 
and is produced by the simultaneous action of the oxygen and the 
water. Ferric hydrate produced in this way can unite chemically 
with ammonia; hence it follows that iron-rust which has been long 
exposed to the air gives off ammonia when heated. 

Many metals which, like zinc, rust in the air, are protected from 
further oxidation by the compact layer of oxide which is first pro- 


Iron. 


443 


duced. But in the case of iron the layer of rust is less compact, 
and the process goes on, comparatively quickly, until the whole of 
ths metal is oxidized. For this reason sheet iron cannot be used to 
protect the roofs of houses, but zinc or lead must be employed 
instead. Iron may be protected from rusting by coating it with a 
layer of some metal which oxidizes slowly in the air, such as tin 
( tin-plate ), zinc {galvanized iron), or lead {terne-plate). The process 
in all cases consists in dipping the clean iron plates into the molten 
metal, and has nothing to do with electricity. Recently a process 
has been introduced of rendering iron articles proof against rust 
by coating them with a layer of the magnetic oxide. For this pur¬ 
pose, the articles are heated to redness in a current of steam, a por¬ 
tion of which oxidizes the surface of the iron while hydrogen is set 
free. It has also been found that iron does not rust in water which 
has been made feebly alkaline by sodium carbonate. 

Iron is the most valuable of all metals ; it is more valuable than 
gold, because its properties in its varied forms as cast-iron, wrought - 
iron, and steel, make it applicable to so many useful purposes for 
which no other metal can be used. And it has been well said that 
without iron there would be no civilization. The metal would be 
still more valuable if some simple means were known by which it 
could be converted into a passive state, so that, like platinum, it 
should remain unchanged in the air and in water containing air 
Such a discovery would completely revolutionize many branches of 
industry. We can, in fact, actually produce this change in the 
properties of the metal. A rod of iron dipped into red fuming 
nitric acid becomes thus passive , and now no longer dissolves in 
ordinary nitric acid, nor rusts in the air. But if touched or scratched 
with a piece of ordinary iron it loses these peculiar properties at 
once. The problem of how to make this passive state permanent 
remains as yet unsolved. 

At a red-heat iron is more rapidly oxidized than by moist air at 
the ordinary temperature. A red-hot rod of iron becomes covered 
with a coating of oxide, which falls off when the iron is hammered 
on the anvil by the smith. This scale , as it is called, consists 
essentially of the magnetic oxide : Fe 3 0 4 , and forms black, dull 
scales when cold. The particles of iron which fly off and burn as 
sparks when a flint and steel are struck together, also form the 
same oxide, and the fused nodules produced when steel burns in 
oxygen (p. 14) have the same composition. 


444 Text-Book of Inorganic Chemistry. 

Red-hot iron absorbs considerable quantities of carbonic oxide, 
which it parts with to the air when in contact with it. This 
happens, for example, when iron stoves are heated to redness. 
And the injurious action of such stoves when too strongly heated is 
largely due to the carbonic oxide which then diffuses into the air 
of the room. 

The iron filings produced by the mechanic in working up iron 
articles are largely used for a number of purposes, owing to the 
larger surface which they expose to the action of various liquids 
than larger pieces of the metal. Many compounds of iron are 
prepared from them, and considerable quantities are employed for 
the commercial reduction of nitrobenzol to aniline in the presence 
of acetic acid. Iron in the state of a fine powder (reduced iro?i) is 
produced, as previously stated, by reducing ferric oxide in a stream 
of hydrogen. 

Iron is attacked and dissolved by nearly all acids; dilute 
sulphuric, hydrochloric, or acetic acid dissolve it, with evolution of 
hydrogen, producing ferrous sulphate, chloride, or acetate respec¬ 
tively. Ordinary concentrated nitric acid dissolves iron, with the 
formation of ferric nitrate and lower oxides of nitrogen. Very 
dilute nitric acid dissolves it, like sulphuric acid; the hydrogen 
which is produced by the decomposition of the water is not, how¬ 
ever, set free, but in its nascent state gradually reduces some of the 
nitiic acid to ammonia. That the solution contains ammonium 
nitrate as well as ferrous nitrate, may be readily proved by the 
ammonia evolved when it is mixed with caustic soda. 

Iron occurs in nearly all its compounds as a dyad and triad 
(or tetrad) element. It is a dyad in the ferrous compounds , ferrous 
oxide : FeO, ferrous chloride : FeCl 2 , &c., and a triad (or tetrad) 
n the ferric compounds — ferric oxide : Fe 2 O a , ferric chloride : 
Fe 2 Cl 6 , &c. Ferrous oxide and ferric oxide both possess basic 
properties , the latter, like all sesquioxides, is a weak base—weaker 
than ferrous oxide. It is, however, more stable than ferrous oxide, 
as it is not changed by the oxygen of the air. A third oxide of 
iron is also known, containing more oxygen than ferric oxide, and 
having the properties of an acid. This ferric acid\ probably 
Fe 0 2 ( 0 H) 2 , appears only to exist in union with strong bases, and 
probably contains hexad iron, like sulphuric acid contains hexad 
sulphur. 


Ferrous Oxide. 


445 


Perrous Oxide : FeO.—On account of the strong attraction of 
this compound for oxygen it is difficult to prepare, and has never 
been obtained pure. It is produced, with some metallic iron, when 
ferrous oxalate is gently heated without access of air, or as a 
black, glistening powder when the mixture of carbonic oxide and 
carbonic acid obtained by heating oxalic acid with sulphuric acid 
(p. 295) is led over hot ferric oxide. The substance so obtained 
easily dissolves in hydrochloric or nitric acid, without evolving 
hydrogen, and when heated in the air is first oxidized to the mag¬ 
netic oxide, and then to ferric oxide. 

Ferrous Hydrate : Fe(OH) 2 , is obtained as a white precipitate 
on mixing solutions of ferrous sulphate and caustic soda both free 
from air. In consequence of the readiness with which it attracts 
oxygen from the air, the white precipitate rapidly becomes green, 
and finally yields a brown mass of ferric hydrate. Ferrous hydrate 
dissolves in salts of ammonium, and is, therefore, not precipitated 
by ammonia from a solution to which ammonium chloride has been 
added. 

Ferrous Sulphate, Green Vitriol : S 0 2 -Q 2 Fe + 7 H 2 0 .—This, 
the best known salt of iron, and that from which most of the other 
salts are prepared, is obtained by dissolving iron in dilute sulphuric 
acid :— 

Fe + S 0 2 ( 0 H) 2 = S 0 2 - 0 2 Fe + H 2 , 

or, as a bye-product in making sulphuretted hydrogen, when ferrous 
sulphide is dissolved in the same acid :— 

FeS + S 0 2 ( 0 H) 2 = S 0 ,- 0 2 Fe + H 2 S. 

It is also produced by gently roasting iron pyrites in the air, 
when a portion of the sulphur burns to sulphurous acid, and another 
portion is oxidized to sulphuric acid, which unites with the iron to 
form ferrous sulphate. Some varieties of pyrites— e.g. marcasite 

_are oxidized at the ordinary temperature when moistened with 

water and exposed to the air, and the decomposed mass yields a 
green solution of ferrous sulphate when treated with water. 

The salt separates from its solution in bluish green crystals with 
7 molecules of water; it is easily soluble in water, especially hot. 
It readily loses 6 molecules of water when heated, but the seventh 
molecule, the water of constitution (p. 86), is only expelled at about 
300°, with a partial decomposition of the salt. The anhydrous salt 


446 Text-Book of Inorganic Chemistry. 

is white, but it regains its green colour when treated with water. If 
heated to redness, the anhydrous salt decomposes into ferric oxide 
and sulphurous and sulphuric anhydrides (p. 164). 

Both solid ferrous sulphate and its solution are partly oxidized 
when exposed to the air, with the formation of basic ferric sulphate 
of a reddish colour ; and when crystals of the substance which 
have been exposed to the air are dissolved in water, this insoluble 
basic salt remains undissolved :— 

6 S 0 2 - 0 2 Fe" + 3O + 3H 2 0 = 2(S0 2 ) 3 0 tf Fe 2 '' + 2Fe'"(OH) 3 . 

Ferrous sulphate unites with the sulphates of the alkalies 
(potassium and ammonium), and forms double sulphates containing 
6 molecules of water— eg. S 0 2 - 0 2 Fe + S 0 2 (OK) 2 + 6aq., corre¬ 
sponding to the magnesium compound (p. 415). These double 
sulphates oxidize less readily than ferrous sulphate itself. The 
salt is an important article of commerce, and is employed for the 
preparation of ferric acetate (used in dyeing), to produce a black 
colour with galls (writing ink), a blue colour with potassium ferro- 
cyanide (Prussian blue), and in the preparation of indigo solutions. 
It is also used as a disinfectant, especially to preserve wood from 
decay. 

Ferrous Carbonate : C 0 - 0 2 Fe.—This compound is found in 
nature as spathic iron ore or siderite , sometimes crystallized in 
rhombohedra which are isomorphous with calcite, but usually in 
compact amorphous masses. Mixed with clay it forms clay iron¬ 
stone— one of the most important of the English ores of iron. The 
salt is obtained as a voluminous white precipitate on mixing a 
solution of ferrous sulphate with one of sodium carbonate. This 
white precipitate gradually absorbs oxygen from the air and changes 
colour, becoming first of a dirty green and then of a brown colour, 
and is finally completely converted into ferric hydrate, while car¬ 
bonic acid is set free—no carbonate containing iron in the ferric 
state is known. If the moist precipitate is rubbed up with sugar, 
it is somewhat protected from the action of the air, and a brownish 
mass is obtained which is used in medicine under the name of ferri 
carbonas saccharata. 

Like calcium carbonate and other carbonates, ferrous carbonate, 
though insoluble in pure water, dissolves in water containing car¬ 
bonic acid. Mineral springs containing ferrous carbonate dissolved 
in carbonic acid are called chalybeate springs. The water has a 
bitter, inky taste, and when exposed to the air becomes turbid from 


Ferric Oxide. 


44 7 


ferric hydrate. The dissolved ferrous carbonate is decomposed by 
the oxygen of the air into ferric hydrate and carbonic acid :— 

2CO 0 . 2 Fe" + O + 3 H 2 0 = 2Fe"'(OH) 3 + 2 C 0 2 . 

PO1 

Ferrous Phosphate : pQ - 0 6 Fe 3 + 8 H-. 0 , is found in nature as 

the mineral vivianite , sometimes crystalline, but usually as an 
earthy powder of a dull blue colour. The compound is precipitated 
when a solution of sodium phosphate is added to one of ferrous 
sulphate, and is used in medicine. It is at first white, but becomes 
blue or green by partial oxidation. By heating iron with phosphoric 
acid, a solution is obtained which after some time deposits colour¬ 
less needles of the composition : PO |q|j C + H 2 0 . 


Ferric Oxide : Fe, 0 3 .—This compound occurs in immense 
quantities in the mineral kingdom, either in lustrous crystals 
belonging to the hexagonal system— specular iron , or more com¬ 
monly in amorphous masses, with or without water, as the important 
iron ores included under the name of hcematite. It may be ob¬ 
tained artificially in the crystalline state by strongly glowing a 
mixture of dry ferrous sulphate and common salt, or in the amor¬ 
phous state by heating dry ferrous sulphate or ferric hydrate. The 
red amorphous powder obtained by either of the last-named 
methods is used for polishing and as a pigment, and is known as 
colcothar , caput mortuum, or rouge. A fine variety of rouge is 
obtained by glowing ferrous oxalate in the ai:. Ferric oxide is 
completely insoluble in water, and after heating to redness only 
dissolves in acids with difficulty. 

Ferric Hydrate : Fe(OH) 3 , is thrown down as a red-brown 
flocculent precipitate when ammonia is added to a solution of a 
ferric salt. When washed with water and dried, it forms a brown 
amorphous mass, insoluble in water. Some important ores of 
iron, and the rust which forms when iron is exposed to the air, con¬ 
sist essentially of ferric hydrate. 

Freshly precipitated ferric hydrate dissolves not only in acids, 
but also in a solution of ferric chloride, producing a dark red solu¬ 
tion. If this solution of basic chloride is placed in a vessel of 
which the bottom is closed with a piece of parchment paper and 
then floated on water, hydrochloric acid diffuses through, and there 
remains in the dialyzer a solution of colloid ferric hydrate as a 





448 Text-Book of Inorganic Chemistry. 


blood-red liquid (dialyzed iron). The addition of a small quantity 
of sulphuric acid, or of an alkali or salt, at once causes the precipi¬ 
tation of the ferric hydrate from this solution as a gelatinous mass. 

The use of ferric hydrate as an antidote in cases of poisoning by 
arsenic has been already referred to (p. 236). 

Like alumina, ferric oxide is a weak base, and its soluble salts 
can only be crystallized with difficulty. The normal salts are white 
or yellowish when anhydrous, red when hydrated ; the basic salts 
are yellow or red. 

Ferric Sulphate : (S 0 2 ) 3 0 6 Fe 2 .—This salt is produced by 
oxidizing a solution of ferrous sulphate with nitric acid, with the 
addition of the one molecule of sulphuric acid which one molecule 
of ferric sulphate contains more than two molecules of ferrous 
sulphate :— 


2S0 2 -0 2 Fe" + S 0 2 ( 0 H) 2 + O = (S 0 2 ) 3 0 6 Fe 2 "' + H 2 0 . 


When the nitric acid is dropped into the hot solution of ferrous 
sulphate and sulphuric acid, the liquid first becomes dark brown 
and nearly black from the compound of ferrous sulphate with the 
nitric oxide produced by the reduction of the nitric acid (p. 194), 
until, when sufficient nitric acid has been added, the whole of the 
nitric oxide is evolved with effervescence, and the dark-red liquid 
now contains ferric sulphate. If no sulphuric acid is added to the 
ferrous sulphate, a larger quantity of nitric acid is required, and 
the solution then contains a mixture of ferric sulphate and ferric 
nitrate. 

On evaporating the aqueous solution of the salt, ferric sulphate 
remains as a white amorphous powder, which attracts water from 
the air, and deliquesces to a dark red liquid, with a bitter astringent 
taste. An excess of ammonia when added to its aqueous solution 
precipitates ferric hydrate, but if less ammonia is added than is 
required to completely decompose it, a brown precipitate of a basic 
salt is produced. The oxidation of ferrous sulphate in the air 
produces a similar compound (p. 446). 

A solution of ferric sulphate, when mixed with one of the potas¬ 
sium sulphates and the mixed solution allowed to evaporate slowly, 
deposits yellowish crystals oi potash-iron alum. The correspond¬ 


ing ammonium compound—ammonia-iron 



449 


Triferric Tetroxide . 

I 2 H 2 0 , crystallizes more readily. The crystals are usually of a 
violet colour, possibly due to organic compounds contained in the 
ammonia. 

Ferric Nit* ate \ (N 0 2 ) 3 0 3 Fe, is prepared by dissolving iron 
in moderately concentrated nitric acid. From the red solution so 
obtained, the salt can only be crystallized with difficulty. The 
solid compound deliquesces in the air. 

Ferric Phosphate (Orthophosphate ) : (PO)O a Fe, is obtained 
as a yellow amorphous precipitate when ferric chloride is mixed 
with sodium phosphate. It is insoluble in water and acetic acid, 
but dissolves in strong inorganic acids. 

Ferric Pyrophosphate closely resembles the preceding salt, and 
is insoluble both in water and acetic acid, but forms a soluble com¬ 
pound with sodium phosphate or pyrophosphate. The soluble 
double pyrophosphate is prepared by adding sodium pyrophosphate 
to the ferric salt until it is just dissolved ; from this solution alcohol 
precipitates the double salt. Its aqueous solution is colourless, and 
does not possess the strong astringent taste of the other soluble 
ferric salts. It is used in medicine. 

Ferric carbonate does not appear to exist. Ferric oxide is toe 
weak a base to unite with carbonic acid. On adding a solution of 
sodium carbonate to one of ferric chloride, a brown precipitate of 
ferric hydrate is produced, and the carbonic acid which is set free 
partly unites with the excess of sodium carbonate to form the acid 
carbonate. 

Triferric Tetroxide, Magnetic Oxide of Iro 7 i\ Fe 3 0 4 = 
Fe^O Fe/^Og.—This oxide is found in nature as black, well- 
developed, regular octahedra, especially in chlorite rocks. It further 
occurs as crystalline masses, and more largely in rough granular 
deposits. Magnetic oxide of iron, so called because pieces of it 
sometimes form natural magnets, is one of the most valuable ores 
of iron. Small lustrous octahedra of the oxide may be obtained 
artificially by heating bright iron wire to redness in a porcelain 
tube, and leading a current of steam over it. The water is then 
decomposed according to the equation 

3Fe + 4H 2 0 = Fe 3 0 4 + 4H 2 . 


G G 


450 Text-Book of Inorganic Chemistry. 

Magnetic oxide of iron is soluble in hydrochloric acid, but the 
solution does not then contain a chloride correspqnding to the 
oxide, but a mixture of ferric and ferrous chlorides. If ammonia 
is gradually added to this solution, a precipitate of ferric hydrate 
is first formed, and then one of ferrous hydrate. But if the solu¬ 
tion is poured, with stirring, into an excess of ammonia, a dark- 
green precipitate is produced, which becomes black and granular 
when boiled. This is the hydrate of the magnetic oxide, and re¬ 
mains unchanged in the air. The same compound may also be 
obtained by converting two-thirds of a solution of ferrous chloride 
into ferric chloride by passing chlorine, then mixing this, when 
the excess of chlorine has been driven off, with the other third of 
the solution of ferrous chloride, and finally pouring the mixed solu¬ 
tion into an excess of ammonia and boiling. 

Ferric Acid. —The free acid, which would have the composi¬ 
tion : Fe vi 0 2 (OH) 2 , has not yet been prepared ; its potassium salt, 
potassium ferrate'. Fe 0 2 ( 0 K) 2 , is obtained by heating a mixture 
of reduced iron with twice its weight of potassium nitrate. The 
oxidation of the iron takes place with a considerable evolution of 
heat, and on extracting the cold mass with water the potassium 
ferrate dissolves with a red colour. On evaporating this solution 
in a vacuum, dark-red crystals are deposited. The salt may also 
be obtained by passing chlorine through concentrated cooled 
caustic potash containing freshly precipitated ferric hydrate in 
suspension :— 

Fe(OH) 3 + 5KOH + 3CI = Fe 0 2 ( 0 K) 2 + 3KCI + 4 H 2 0 . 

Iron forms two compounds with each of the halogens—chlorine, 
bromine, and iodine—corresponding to ferrous and ferric oxide 
respectively. 

Ferrous Chloride : FeCl 2 .—The anhydrous compound is ob¬ 
tained by heating iron wire in a stream of hydrochloric acid gas, 
when it sublimes at a high temperature in white, glistening iri¬ 
descent crystals which feel soft and talc-like to the touch. It 
melts at a red-heat, and then solidifies to a crystalline mass. It is 
easily soluble in water, and on evaporating the aqueous solution 
without access of air, green crystals of the compound : FeCl 2 + 
4 H 2 0 separate out. The same compound may be obtained by 
dissolving iron in hydrochloric acid and evaporating down. A 
solution of ferrous chloride, when exposed to the air, is gradually 
oxidized to basic ferric chloride. 


Ferric Chloride . 


451 




Ferric Chloride : Fe 2 Cl 6 .—Although ferrous chloride readily 
unites with a further quantity of chlorine, an atom of iron can 
never decompose more than two molecules of hydrochloric acid. 
And when iron is dissolved in the acid, ferrous chloride is always 
produced, never ferric chloride. Anhydrous ferric chloride may 
be obtained by heating iron in dry chlorine, and then sublimes in 
nearly black, lustrous, crystalline plates. It is far more volatile, 
and therefore sublimes much more easily than ferrous chloride. 
When exposed to the air it becomes yellow, and deliquesces, 
forming a yellow or reddish solution ; it is also soluble in alcohol 
and ether, and is even abstracted from its aqueous solution when 
shaken with ether. The aqueous solution, which is also obtained 
by dissolving iron in aqua regia, or by passing chlorine through a 
solution of ferrous chloride, deposits reddish, deliquescent crystals 
of the compound: Fe 2 Cl 6 + I2H 2 0, when evaporated to a syrup. 
This compound, when heated, decomposes into water and hydro¬ 
chloric acid, which pass off with some anhydrous ferric chloride, 
while ferric oxide remains behind. An aqueous solution of ferric 
chloride can dissolve freshly precipitated ferric hydrate in con¬ 
siderable quantities, producing a dark-red liquid, from which ferric 
hydrate is again precipitated on diluting with water and heating. 

Ferric chloride unites with the chlorides of potassium and 
ammonium to form double salts, which are deposited in bright-red 
crystals of the composition: Fe 2 Cl 6 4 4 KCI + 2H 2 0 and Fe 2 Cl ti 4 
4 NH 4 C 1 + 2 H 2 0 , when the mixed solutions are evaporated down. 

The bromides of iron closely resemble the chlorides. 

Ferrous Iodide : Fel 2 .—This compound is obtained by heating 
iron filings with iodine in a porcelain crucible. A small quantity 
of iodine is first added, then an excess, and the heating continued 
as long as vapours of iodine are evolved. The salt then remains 
as a dark green lamellar mass, easily soluble in water. An 
aqueous solution of ferrous iodide may also be obtained by digest¬ 
ing iron filings with water and iodine—the iron being in excess. 
On evaporating this solution, crystals of the composition : Fel 2 + 
4 H O, are deposited. This solution dissolves more iodine when 
digested with it, and the dark-red solution so produced probably 
contains ferric iodide : Fe 2 I 6 , which has not yet been obtained in 
the solid form. 

The compounds of iron with sulphur do not exactly correspond 

g g 2 


452 Text-Book of Inorganic Chemistry. 

with its oxides. A sesquisulphide of the composition : Fe.^S 3 , and 
corresponding to ferric oxide, is not yet known for certain. On 
the other hand, the sulphide most widely distributed in nature is 
the disulphide : FeS 2 , to which no corresponding oxide is known. 

Ferrous Sulphide : FeS, is prepared as a black crystalline 
mass, by heating three parts of iron filings with nearly two parts 
of flowers of sulphur. If an excess of sulphur is employed, no 
compound richer in sulphur is obtained, but the excess is volatilized. 
Iron and sulphur unite together even at ordinary temperatures 
if the finely powdered mixture is moistened with water. When 
larger quantities are used the heat of combination is often suffi¬ 
cient to violently expel a portion of the mixture from the vessel 
in which the reaction takes place. 

Ferrous sulphide is insoluble in water, but is decomposed by 
acids. With hydrochloric or dilute sulphuric acid it yields con¬ 
siderable quantities of sulphuretted hydrogen, and is therefore 
largely used to prepare this gas. 

Ferric Disulphide : FeS.,.—This compound is widely distri¬ 
buted in the mineral kingdom, either in bright yellow crystalline 
masses, or in well-developed lustrous cubes, and other forms 
belonging to the regular system. In these forms it is known as 
iron-ftyrites. It may also be prepared in the form of a yellow 
powder by gently heating powdered ferrous sulphide with about its 
own weight of sulphur, until the excess of this latter substance is 
expelled. Iron-pyrites is so hard that it strikes sparks with steel ; 
unlike ferrous sulphide, it is not attacked by dilute acids, and is not 
therefore adapted for the preparation of sulphuretted hydrogen. 
When heated strongly it parts with some of its sulphur, leaving 
compounds of ferrous sulphide with more or less sulphur, accord¬ 
ing to the* degree and duration of the heating. If iron-pyrites 
is heated in the air, a large portion of the sulphur is oxidized to 
sulphurous anhydride, and for this reason immense quantities of 
pyrites are employed in the manufacture of sulphuric acid. Many 
kinds of iron-pyrites contain small quantities of copper-pyrites, 
and traces of the noble metals, and the residues of these pyrites 
after roasting are often worked for copper and the other metals. 

Besides iron-pyrites, which crystallizes in the regular system, a 
second modification of ferric disulphide is known, the crystals of 
which belong to the rhombic system, and which is distinguished 


453 


Compounds of Iron and Cyanogen. 

from iron-pyrites by the readiness and rapidity with which it oxidizes 
when exposed to moist air—evolving at the same time a consider¬ 
able amount of heat. This variety of pyrites is called marcasite ; 
and the chief product of its oxidation is ferrous sulphate, which is 
obtained in considerable quantities from it. 

A third sulphide of iron is also found in nature, which contains 
less sulphur than iron-pyrites, and not unfrequently has the com¬ 
position : Fe 7 S 8 , or, if considered as a compound of ferrous sul¬ 
phide and iron-pyrites : 6FeS,FeS 2 . This bronze-coloured ipineral 
is known as magnetic pyrites , because it is attracted by the magnet. 
It crystallizes in the hexagonal system. 


COMPOUNDS OF IRON AND CYANOGEN. 

Although the compounds of iron with cyanogen corresponding 
to ferrous and ferric chlorides have not yet been prepared pure, 
the double salts which ferrous and ferric cyanide yield with the 
cyanides of other metals are among the best known and most 
stable of chemical compounds. Various views are held as to the 
chemical constitution of these substances, but in the present state 
of our knowledge it is perhaps best to write their formula em¬ 
pirically, without entering into the question of their constitution. 
The best known of these, and that from which all the others are 
prepared, is 

Potassium Ferrocyanide : K 4 Fe // Cy 6 + 3^0 = 4ECy, Fe Cy 2 

+ 3 H 2 0 ._This compound, known in trade as yellow prussiate of 

potash , or simply as yellow prussiate , crystallizes from its aqueous 
solution in large, pale-yellow, quadratic plates, which are soft, 
and therefore difficult to powder. Heated to ioo°or a little above, 
it loses its water of crystallization, and then becomes white and 
brittle, so that it can be easily pulverized. It is soluble in water, 
but insoluble in alcohol; ioo parts of water dissolve 25 parts at 
the ordinary temperature and 50 parts at the boiling-point. 

Potassium ferrocyanide is obtained from potassium cyanide—a 
salt which is extremely poisonous, and which is readily decom¬ 
posed by dilute acids. From this it might be expected that the 
ferrocyanide would also be poisonous, but experiment has proved 
that considerable quantities may be taken internally without any in¬ 
jurious or even unpleasant effects. When taken into the system it 




454 Text-Book of Inorganic Chemistry . 

soon gets into the blood, and in a short time can be detected in the 
urine, in which it is expelled unchanged from the body. 

The salt is readily obtained from potassium cyanide by digest¬ 
ing an aqueous solution with iron filings or powdered ferrous 
sulphide. In the former case, caustic potash and hydrogen are 
also produced ; in the latter potassium sulphide :— 

6KCy + Fe + 2 H 2 0 = K 4 FeCy 6 + 2KOH + H 2 

and 

6KCy + FeS = K 4 FeCy 6 + K 2 S. 

The same process goes on in the manufacture of potassium 
ferrocyanide on a large scale, except that instead of using potassium 
cyanide itself materials are employed which, when heated with 
potassium carbonate, produce this substance. For this purpose, 
organic compounds rich in nitrogen, such as blood, horn, hair, 
leather, are mixed with potash and iron filings, and fused at a red 
heat. During the process the carbon and nitrogen of the organic 
substances unite with the potassium and form potassium cyanide ; 
and when the fused mass is afterwards extracted with water this 
substance acts upon the iron, which has been partly converted into 
ferrous sulphide by the sulphur of the crude materials, in the 
manner explained above. The salt which separates out on evapo¬ 
ration is purified by recrystallization ; it often contains potassium 
sulphate. 

Potassium ferrocyanide is not acted upon by carbonic acid nor 
even by dilute acids, and the potassium cyanide in combination 
with the ferrous cyanide has, therefore, acquired other properties, 
and has become more stable. The iron, too, has lost many of the 
properties which are possessed by the ferrous salts ; ammonia does 
not precipitate ferrous hydrate, nor does ammonium sulphide 
throw down ferrous sulphide. And that potassium ferrocyanide is 
not an ordinary double salt of 4 molecules of potassium cyanide 
and 1 molecule of ferrous cyanide is proved by the fact that it is 
not decomposed by dilute acids like other true double cyanides. 
The double cyanide of nickel and potassium : 2KCy,NiCy 2 , a 
compound soluble in water, at once gives a precipitate of nickel 
cyanide with dilute hydrochloric acid, while the potassium cyanide 
is decomposed into potassium chloride and hydrocyanic acid. But 
no ferrous cyanide is separated by hydrochloric acid from the 
apparently analogous compound potassium ferrocyanide. 


Potassiu in Ferro cyanide. 455 

Many attempts have been made to explain this peculiar be¬ 
haviour of potassium ferrocyanides and other similar compounds, 
but as yet without success. That potassium ferrocyanide when 
heated with sulphuric acid leaves a compound containing I mole¬ 
cule of ferrous cyanide and i molecule of potassium cyanide, 
seems to prove that one of the four molecules of potassium 
cyanide in the ferrocyanide is more firmly combined than the other 
three. 

An aqueous solution of potassium ferrocyanide is decomposed 
by strong acids (hydrochloric or sulphuric acid), with the forma¬ 
tion of hydroferrocyanic acid : H 4 FeCy 6 , and the potassium salt of 
the acid employed. If finely powdered potassium ferrocyanide is 
heated with dilute sulphuric acid, hydrocyanic acid is evolved (p. 313), 
and there remains potassium sulphate with the compound we have 
just mentioned : KCy,FeCy 2 . This is a white substance, turning 
blue in the air. Potassium ferrocyanide, when powdered and heated 
with concentrated sulphuric acid, undergoes a further decomposi¬ 
tion. There are produced ferrous, potassium, and ammonium 
sulphates with the evolution of a gas which consists chiefly of car¬ 
bonic oxide. But since traces of other gases are mixed with the 
carbonic oxide, this method is not suitable for preparing the pure 
substance. 

When heated in an iron crucible, potassium ferrocyanide 
is decomposed into iron carbide and potassium cyanide (p. 359 )- 
Fused with manganese peroxide or red-lead, potassium cyanate 
and ferric oxide are produced (p. 361). And when mixed with 
potassium carbonate and sulphur and fused, potassium sulpho- 
cyanate results (p. 361). If chlorine is led through a solution of 
potassium ferrocyanide, the iron is oxidized to the ferric state and 
some of the potassium converted into chloride, producing a com¬ 
pound called potassium ferricyanide : K 3 Fe /// Cy 6 = 3 KCy,Fe w Cy 3 

(P Potassium ferrocyanide, like all other soluble ferrocyanides, is 
distinguished by its reactions with ferric chloride and soluble 
copper salts. When highly diluted it gives with the former a dark 
blue precipitate of Prussian blue, and with the latter a chocolate 
brown precipitate of copper ferrocyanide. With a soluble lead 
salt a solution of potassium ferrocyanide gives a white precipitate 
of the corresponding lead salt. The following equations represent 
these reactions :— 


456 


Text-Book of Inorganic Chemistry. 

Up'// ) 

2K 4 Fe"Cy 6 + Fe 2 "'Cl 6 - 6 KC 1 + 2^ e JCy 4 ,Fe"Cy 2 . 

Prussian blue 

K 4 FeCy 6 + 2 CuC 1 2 = 4KCI + Cu 2 FeCy 6 . 

Copper 

ferrocyanide 

K 4 F eCy, + 2 (N 0 ,), 0 4 Pb - 4N0 2 -0K + Pb 2 FeCy c . 

Lead 

ferrocyanide 

The majority of the ferrocyanides are insoluble in water; those 
of the alkalies and alkaline earths are soluble, but that of barium 
only with difficulty. 

Sodium Ferrocyanide : Na 4 FeCy 6 + I 2 H 2 0 .—This salt cor¬ 
responds closely to the potassium compound, except that it contains 
considerably more water of crystallization. It may be obtained 
from a Prussian blue which contains an atom of hydrogen in 
place of potassium in the usual compound, and which is pre¬ 
cipitated from an acid solution. If this Prussian blue is boiled 
with caustic soda, ferric hydrate and soluble sodium ferrocyanide 
are produced :— 

^ e jcy 4 ,FeCy 2 + 4NaOH = Fe(OH) s + Na 4 FeCy 6 + H 2 0 . 

Other soluble ferrocyanides may be obtained in the same way. 

Ammonium Ferrocyanide : NH 4 FeCy 6 + 3 H 2 0 , crystallizes in 
bright yellow quadratic prisms, and is isomorphous with the potas¬ 
sium salt. 


Fe w ) 

Prussian Blue : rCy 4 ,FeCy 2 . 

Various dark blue compounds are known under the name of 
Prussian blue, but the commonest insoluble compound has the 
above composition ; it is obtained on adding potassium ferro¬ 
cyanide to a solution of ferric chloride or some other ferric salt. 
The compound may be considered as potassium ferrocyanide in 
which three atoms of potassium have been displaced by one of 
triad iron. The atom of potassium which it contains may be dis¬ 
placed by hydrogen without any change in the physical properties 
of the substance. Such a compound may be obtained by precipi¬ 
tating hydroferrocyanic acid with ferric chloride, or by adding 


Prussian Blue. 457 

potassium ferrocyanide to a solution of ferric chloride containing 
free hydrochloric acid. 

Prussian blue can only be washed on the filter with difficulty ; 
when dry it forms a deep blue, hygroscopic powder, which under 
pressure assumes a lustre like copper. It is slightly soluble in 
dilute acids, and when warmed with concentrated acids is decom¬ 
posed. Oxalic acid dissolves the compound when free from 
potassium, and gives a deep blue liquid, which is used as a blue 
ink. Alkalies and alkaline carbonates and even the alkalies of 
soap decompose it, separating ferric hydrate and reproducing 
potassium ferrocyanide :— 

£ e }c y4) FeCy, + 3 KOH = K 4 FeCy 6 + Fe(OH) 3 . 

Prussian blue cannot, therefore, be used for dyeing or printing 
articles which are to be washed with soap. 

A variety of Prussian blue, soluble in water to a deep blue 
solution, is obtained when ferric chloride is added to a large 
excess of potassium ferrocyanide. This soluble Prussian blue is 
probably a compound of ordinary Prussian blue and potassium 
ferrocyanide. 

Hydroferrocyanic Acid : H 4 FeCy 6 .—This acid compound is 
obtained as small white crystalline plates when a concentrated 
solution of potassium ferrocyanide free from air is mixed with 
strong hydrochloric acid in the cold :— 

K 4 FeCy 6 + 4HCI = 4KCI + H 4 FeCy 6 . 

The crystals separate more readily if a little ether is added. They 
are afterwards brought upon a filter and washed with water con¬ 
taining ether. The compound is soluble in alcohol and in water, 
but is insoluble in hydrochloric acid and ether. It behaves like a 
tetrabasic acid, and when mixed with potassium carbonate again 
produces potassium ferrocyanide. It is rapidly oxidized in the 
air and becomes of a blue colour. 

Potassium Perricyanide : KgFe'^Cyg = 3KCy,Fe // 'Cy 3 .— 
This compound, which is known in trade as red prussiate , differs 
from the ferrocyanide by containing its iron in the ferric state. It 
is produced by abstracting potassium from the ferrocyanide, when 
the cyanogen thus set free unites with the iron and changes it 
from the ferrous to the ferric state. This is best done by leading 


458 Text-Book of Inorganic Chemistry. 

chlorine into a solution of potassium ferrocyanide until a portion 
of the liquid no longer gives a blue precipitate with ferric chlo¬ 
ride :— 

K 4 FeCy 6 + Cl = KC 1 + K 3 FeCy 6 . 

The salt separates out on evaporating its aqueous solution in 
dark-red, lustrous, rhombic prisms, easily soluble in water and in¬ 
soluble in alcohol. Its aqueous solution is of a dirty green colour 
and decomposes on standing, especially when exposed to sunlight, 
depositing a blue precipitate. The readiness with which it again 
forms potassium ferrocyanide, especially in the presence of alka¬ 
lies, make it a powerful oxidizing agent. 

Potassium ferricyanide gives no blue precipitate with ferric salts, 
but only a brown colour, but with ferrous salts it gives a fine blue 
precipitate, closely resembling Prussian blue, and called— 

Fe") 

Turnbull’s Blue : ^ h Cyg^e^Cyg. — The precipitate is 
formed according to the equation :— 

3 KCy,FeCy 3 + FeCl ; = 2 KC 1 + £ e j Cy 3 ,FeCy 3 . 

The compound, like Prussian blue, is decomposed by caustic 
alkalies, reproducing potassium ferricyanide, while ferrous hydrate 
is precipitated. At the same time, and especially with excess of 
alkali, a portion of the ferrous hydrate is oxidized to ferric hydrate 
by the potassium ferricyanide, which is reduced to the ferro¬ 
cyanide. 

Hydroferricyanic Acid : H 3 FeCy 6 .—This compound separates, 
like the corresponding hydroferrocyanic acid, on mixing a solution 
of potassium ferricyanide with strong hydrochloric acid in the cold. 
It forms brown crystals, which are easily decomposed. 


Sodium Nitroprusside : Na 2 Fe /// Cy 5 (NO) + 2 H 2 0 . — This salt, 
which contains one atom of sodium less than the ferricyanide, 
and one atom of cyanogen displaced by the monad radical: NO 
(nitryl), crystallizes more easily than that of potassium. It may 
be prepared in the following manner. Two parts of potassium 
ferrocyanide are gently warmed with 3 parts of strong nitric acid 
which has been diluted with an equal volume of water, until a 
portion of the liquid no longer gives a blue precipitate either with 
ferric chloride or with ferrous sulphate. The liquid is then allowed 


Detection of Iron Compounds. 459 

to cool, so that potassium nitrate which has been produced shall 
crystallize out; the acid liquid is then poured off, diluted with water, 
and neutralized with sodium carbonate. 

On evaporating this solution sodium nitroprusside crystallizes 
out on cooling in large, dark-red, transparent crystals, easily soluble 
in water and in dilute alcohol. The aqueous solution gradually 
decomposes, especially when exposed to light, and deposits a 
blue precipitate. Sodium nitroprusside does not precipitate ferric 
salts, but is especially characterized by the dark purple colour 
which it produces with alkaline sulphides, and even with sulphu¬ 
retted hydrogen. The colour soon changes to a blue, and after¬ 
wards becomes a dirty brown. This reaction is so delicate that 
by its means the presence of the most minute traces of an alkaline 
sulphide or of a nitroprusside may be detected. 


Detection of Iron Compounds. 

The presence of small quantities of iron in any form may be re¬ 
cognized by the colour of the borax bead produced in the blow¬ 
pipe flame. In the outer or oxidizing flame, the ferric oxide colours 
the bead yellow, the colour becoming fainter on cooling. In the 
inner or reducing flame, the bead becomes of an olive-green colour, 
due to ferrous oxide. These colours cannot be observed in the 
presence of some other metals— e.g. cobalt—which obscure the 

colours produced by the oxides of iron. 

An aqueous solution of a ferrous salt gives no precipitate with 
sulphuretted hydrogen, but a black precipitate of ferrous sulphide 
with ammonium sulphide. Caustic soda or ammonia precipitates 
ferrous hydrate from the same solution ; the precipitate is first 
white, but soon oxidizes and becomes green and then brown. 
Potassium ferrocyanide gives a white precipitate rapidly oxidizing 
to Prussian blue, while potassium ferricyanide gives a blue pre¬ 
cipitate (Turnbull’s blue) at once. . 

When a stream of sulphuretted hydrogen is led through a 
solution of a ferric salt, a white precipitate of sulphur falls down, 
and the ferric salt is reduced to the ferrous state 

Fe 2 Cl 6 + H 2 S = 2FeCl 2 + 2HCI + S. 

And the black precipitate which ammonium sulphide produces in 
ferric salts is a mixture of ferrous sulphide and sulphur. Caustic 
soda or ammonia, when added to a solution of a ferric salt, pre- 


460 Text-Book of Inorganic Chemistry . 

cipitates brown ferric hydrate ; the presence of potassium tartrate 
prevents the precipitation. 

If a ferric salt containing phosphoric acid is precipitated with 
ammonia, the precipitate is not ferric hydrate, but ferric phosphate. 
Potassium ferrocyanide when added to a ferric salt gives a dark- 
blue precipitate of Prussian blue. Potassium ferricyanide produces 
no precipitate, but only a brown colour. 

Iron, in the form of a ferric salt, may be separated from the 
metals of the alkaline earths by precipitating with ammonium 
chloride and ammonia. Ammonium sulphide precipitates iron in 
any form, but does not throw down the alkaline earths. 

Iron is separated from aluminium by digesting with an excess 
of caustic soda, which dissolves aluminium hydrate, but not ferric 
hydrate. 


MANGANESE. 

Chemical Symbol'. Mn. —Atomic Weight : 55. 

Manganese is closely related to iron in its properties, and has 
about the same atomic weight. Its compounds are also nearly 
always found associated with iron in nature. The most important 
minerals containing manganese are its oxygen compounds, and of 
these pyrolusite (manganese peroxide : Mn 0 2 ) is found in by 
far the largest quantities. Other manganese minerals are brau- 
nite (manganic oxide : Mn 2 0 3 ), manganite (manganic hydrate : 

{ MnO -OH^ hausmanmte (trimanganic tetroxide : Mn 3 0 4 ), man- 

ganese-spar (manganous carbonate: CO • 0 2 Mn), psilomelane and 
ivad (impure manganic perhydrate). Manganese also occurs in 
nature as its silicate, and in combination with sulphur as various 
sulphides. 

The metal itself is of but little interest. It may be obtained by 
heating a mixture of precipitated manganous carbonate and oil to 
redness, and then again heating this mixture with charcoal powder 
and anhydrous borax to the highest temperature of a wind furnace. 
The regulus of metal so obtained, which contains carbon and other 
impurities, is of a grey colour with a reddish tinge, very hard and 
brittle, and with a high melting-point. It has a specific gravity of 
about 7*2, and oxidizes when exposed to moist air. The metal 



Manganese. 


461 


dissolves easily in dilute sulphuric, hydrochloric, or even nitric acid, 
producing the corresponding manganous salts. 

Manganese unites with oxygen in no less than five proportions, 
forming oxides of very different chemical properties, and in which 
the manganese plays the part of a dyad, triad, tetrad, hexad, and 
heptad element. These compounds are the following :— 


Manganous oxide : MnO 
Manganic oxide : Mn 2 0 3 
Manganese peroxide : Mn0 2 
Manganic acid : H 2 Mn0 4 
Permanganic acid : HMn0 4 


Strong base 
Very weak base 
Indifferent 
Dibasic acid 
Monobasic acid. 


Of these oxides, manganous oxide corresponds to ferrous oxide, 
manganic oxide to ferric oxide, and an intermediate compound of 
these two : Mn 3 0 4 , to the magnetic oxide of iron. An oxide of iron 
corresponding to manganese peroxide is unknown, but we are ac¬ 
quainted with the corresponding sulphide in iron-pyrites. Finally, 
manganic acid corresponds to ferric acid, but no compound of 
iron is known similar to permanganic acid. 


Manganous Oxide : MnO.—This oxide remains as a greenish 
powder when manganous carbonate is glowed out of contact with 
air, or in a stream of hydrogen. It may also be obtained by heat¬ 
ing a mixture of manganous chloride and sodium carbonate. Man¬ 
ganous oxide is a strong base, and dissolves easily in hydrochloric, 
sulphuric, or nitric acid. When heated in the air it absorbs oxygen 
and is converted into trimanganic tetroxide : Mn 3 0 4 . Unlike the 
iron compound, it is not reduced to the metallic state by hydrogen, 
even at a high temperature. 

Manganous Hydrate'. Mn(OH) 2 , is deposited as a white pre¬ 
cipitate when caustic soda is added to a solution of a manganous 
salt. It rapidly darkens, especially when exposed to the air, and 
finally becomes completely converted into dark brown manganic 
hydrate. 

Ammonia produces the same precipitate when the solution does 
not contain ammonium salts; manganous hydrate forms soluble 
compounds with these substances. 

The most important of the manganous compounds — i.e. those in 
which the metal is a dyad—are the following 

Manganous Sulphate : S0 2 • 0 2 Mn + 7H 2 0.—This salt is pro- 


462 Text-Book of Inorganic Chemistry . 

duced, with evolution of oxygen, when manganese peroxide is heated 
with strong sulphuric acid :— 

Mn0 2 + S0 2 (0H ) 2 = S0 2 -0 2 Mn + H,0 + O. 

The filtered solution, when sufficiently concentrated and freed 
from iron, deposits the salt in clear, rose-coloured, tabular crystals. 
At the ordinary temperature the salt crystallizes with only five 
molecules of water, and at higher temperatures other hydrated salts 
may be obtained. Like magnesium and ferrous sulphates, the salt 
unites with the alkaline sulphates, forming double sulphates with 
six molecules of water. 

Manganous Nitrate : (N0 3 ) 2 -0 2 Mn.—If manganous carbonate 
is dissolved in nitric acid and the solution evaporated down, man¬ 
ganous nitrate is obtained, and the salt is not oxidized as in the 
case of iron. Manganous nitrate is highly deliquescent, and can 
therefore only be crystallized with difficulty ; when dried and heated 
it darkens in colour, and manganese peroxide separates out. 

Manganous Carbonate: C0-0. 2 Mn, is thrown down as a 
white precipitate when sodium carbonate is added to a solution 
of a manganous salt. It remains unchanged in the air when moist 
and only slightly darkens when dried. The salt is therefore much 
more stable than ferrous carbonate. Manganous carbonate occurs 
in nature as the mineral mahganese-spar , crystallizing in rhombo- 
hedra, isomorphous with spathic iron ore and calcite. 

Manganous Chloride : MnCl 2 + 4H 2 0.—Of the two chlorides 
of iron, ferric chloride is the more stable, and ferrous chloride 
readily combines with chlorine. But manganous chloride is by 
far the most stable chloride of the metal. The compound is pro¬ 
duced when manganous carbonate is dissolved in hydrochloric acid, 
and is also a bye-product in the preparation of chlorine from man¬ 
ganese peroxide. As already stated (p. 98 ), manganese tetra¬ 
chloride : MnCl 4 , is first produced, but is so unstable that it at once 
breaks up into manganous chloride and free chlorine. But as the 
manganese peroxide used for the preparation of chlorine always 
contains iron, the manganous chloride produced when it is dis¬ 
solved in hydrochloric acid always contains ferric chloride. 
Various methods are used to separate the iron, one of which 
consists in heating the dried residue in a Hessian crucible to low 
redness : the ferric chloride is then partly volatilized and partly 


Manganous Compounds . 463 

converted into insoluble basic chloride, but the manganous chloride 
remains unchanged. The mass is afterwards powdered and ex¬ 
tracted with hot water; and on evaporating the violet-coloured 
solution, manganous chloride with four molecules of water separates 
out in pink-coloured crystals. It is very easily soluble in water, and 
deliquesces in moist air; when heated, it parts with all its water, 
accompanied by a trace of hydrochloric acid. Manganous chloride 
unites with ammonium chloride and forms a crystalline double 
salt of the composition : MnCl 2 ,2NH 4 Cl + H 2 0. 

Manganous Sulphide : MnS, is found in the mineral kingdom 
as alabandite , in black cubes, or in crystalline masses. It is 
obtained artificially as a flesh-coloured precipitate by mixing a 
solution of a manganous salt with one of sodium sulphide. It easily 
dissolves in dilute acids with evolution of sulphuretted hydrogen. 
—Manganese Disulphide : MnS 2 , has not been prepared artificially, 
but occurs in nature as the mineral hauerite in reddish-brown 
crystals belonging to the regular system. 

Manganic Oxide : Mn 2 0 3 , occurs in nature as the mineral 
braunite in black lustrous crystals belonging to the quadratic 
system. 

Manganic Hydrate'. |JJnO*OH * s *° unc * a * so m natur e as 

manganite in dark-brown rhombic prisms or crystalline masses. 
It is prepared as a dark-brown amorphous precipitate when a 
solution of manganous chloride is mixed with ammonium chloride 
and ammonia and exposed to the air. Unlike manganous hydrate, 
it is insoluble in ammonium chloride. 

Manganic oxide is a weak base, and difficultly combines with 
acids. Its salts decompose easily, especially when warmed, and 
are converted into manganous compounds. 


Manganic Sulphate : (S0 2 ) 3 0 6 Mn 2 /// , may be obtained as a 
dark-green amorphous powder by digesting pure manganese per- 
hydrate with concentrated sulphuric acid. It deliquesces when 
exposed to the air, forming a red liquid, and is a difficult compound 
to prepare. Somewhat more stable is its double compound with 

potassium sulphate, manganese alum : so 2 }oK^ n+ I2 ^2 O, which 
is obtained by mixing the solution of the two sulphates. 


464 Text-Book of Inorganic Chemistry. 

Manganic Chloride is a very unstable compound, and unknown 
in the solid state. When manganic hydrate is treated with hydro¬ 
chloric acid in the cold a dark brown liquid is obtained, which 
probably contains manganic chloride, but which continually evolves 
chlorine. 

Triman-anic Tetroxide : Mn 3 0 4 = Mn0,Mn 2 0 3 , is found in 
nature as the mineral haus?nannite in brown quadratic crystals. 
It is the most stable oxide of manganese, and may be heated to 
redness in the air without change. If manganese peroxide is 
strongly heated as long as oxygen is evolved, this compound 
remains as a reddish powder, and when manganous oxide or car¬ 
bonate is heated in the air it absorbs oxygen and produces the 
same substance. Hydrochloric acid dissolves trimanganic tetroxide, 
forming manganous chloride, while chlorine is evolved. 

Manganese Peroxide : Mn0 2 , is the commonest and most 
important of the manganese minerals, and as such is known by the 
name pyrolusite. It crystallizes in grey rhombic prisms, with a 
metallic lustre, but is usually found in rough masses often fibrous 
in structure. The chief European localities where it is found are 
Thuringia, and on the Lahn, in Germany, and in Bohemia, France, 
and Spain. Manganese peroxide gives a grey streak on paper or 
porcelain, and may thus be distinguished from other manganese 
minerals, which give a brown streak. The compound may be 
artificially prepared by carefully heating manganous carbonate 
with potassium chlorate and extraction with water, or by heating 
manganous nitrate or the hydrated peroxide. 

Manganese peroxide is an indifferent substance, insoluble in 
water and in nitric acid, even when concentrated. Strong sulphuric 
acid dissolves it with formation of manganous sulphate and free 
oxygen ; hydrochloric acid converts it into manganous chloride 
and chlorine. Substances which are easily oxidized, such as sugar, 
oxalic acid, &c., reduce manganese peroxide in the presence of 
sulphuric acid. Oxalic acid is thus completely converted into 
carbonic acid and water, while the peroxide becomes manganous 
sulphate. A process based upon this reaction is used to determine 
the quantity of the peroxide contained in commercial pyrolusite. 
Every molecule of manganese peroxide oxidizes a molecule of 
oxalic acid to carbonic acid, and it is, therefore, only necessary to 
know the weight of the latter substance, which may be either 


465 


Manganese Peroxide. 


collected and weighed, or estimated from the loss in weight, in 
order to determine the value of any particular sample of the ore : 

jcO-OH + Mn ° 3 + S0 2 (0H ) 2 - 2 CO a + S0 2 -0 2 Mn + 2 H 2 0 . 


Manganese peroxide is largely used for many technical purposes, 
chiefly for the manufacture of chlorine and bleaching powder. 
When heated, it liberates considerable quantities of oxygen, and 
becomes converted into trimanganic tetroxide. 


Manganese Perhydrate , of uncertain composition, is obtained as 
a dark-brown amorphous powder when a solution of manganous 
chloride is precipitated by sodium hypochlorite, or when chlorine 
is passed through a solution containing suspended manganous 
hydrate. The hydrate also separates when carbonic acid is passed 
through a solution of potassium manganate, which is then changed 
into potassium permanganate. The same change takes place if a 
very dilute solution of potassium manganate is boiled. 

1 0 H 

Qpj—The acid of the above com¬ 
position has not yet been prepared ; in all attempts to separate it, 
it breaks up into permanganic acid, water, and manganese perhy¬ 
drate. From the composition of its salts we infer that it is a dibasic 
acid, having the above formula. 

{OK 

Potassium Manganate: Mn0 2 |Q K —This salt is obtained by 

the oxidation of any oxide of manganese in the presence of free 
alkali—for example, when any compound of manganese is fused 
with a- mixture of potassium (or sodium) carbonate, and a little 
nitre. It may be obtained purer by the following process. Five 
parts of powdered caustic potash are mixed with a little water and 
4 parts of powdered manganese peroxide stirred up with it; this 
mixture is next heated nearly to boiling, and 3 ^ parts of powdered 
potassium chlorate added. After evaporating the mass to dry¬ 
ness, it is heated in a crucible to low redness for one and a half 
hours, and the cold mass extracted with a little cold water. The 
intensely green solution of potassium manganate so obtained 
deposits dark green, almost black, rhombic crystals of the salt 
when evaporated in a vacuum over sulphuric acid. The reaction 
is represented by the following equation 

3 Mn 0 2 + 6K0H + C10 2 -OK = 3 Mn 0 2 ( 0 K ) 2 + KC1 + 3H 2 0. 

Potassium manganate is an unstable compound ; its aqueous 

H H 


466 Text-Book of Inorganic Chemistry. 

solution changes on standing, especially when warmed, and be¬ 
comes red from formation of potassium permanganate—hence the 
name chameleon mineral sometimes given to the salt. 


Permanganic Acid: Mn0 3 -0H.—This monobasic acid is only 
known in solution, and is so unstable that it rapidly decomposes. 
On the other hand, the anhydride —permanganic anhydride : 


Mn Q 0 7 = ^jno 3 } may be obtained by addin S potassium perman¬ 
ganate in small quantities at a time to well-cooled strong sulphuric 
acid. It then separates from the olive-green liquid as drops of a 
heavy green oil, which soon decompose with evolution of oxygen. 


Potassium Permanganate'. Mn0 3 -0K.—This compound, the 
best known of the permanganates, is obtained by boiling the green 
solution of potassium manganate, or better, by leading a stream of 
carbonic acid through it. In the former case there are produced 
potassium permanganate, manganese perhydrate, and caustic 
potash ; in the latter case, potassium carbonate is formed instead 
of the hydrate :— 

3 Mn 0 2 ( 0 K ) 2 + 2 H 2 0 = 2 Mn 0 3 - 0 K + Mn0 2 + 4 KOH. 
3 Mn 0 2 ( 0 K ) 2 + 2C0 2 = 2 Mn 0 3 - 0 K + Mn0 2 + 2 CO-(OK) 2 . 

For the preparation of the salt a solution of potassium man¬ 
ganate is warmed and a stream of carbonic acid led through it 
until the solution has become of a bright red qolour. The dear 
solution is then run off from the precipitate and evaporated down, 
when potassium permanganate separates on cooling in long, dark- 
red, almost black, needles, with a metallic lustre. The salt is 
iso-morphous with potassium perchlorate. It is tolerably easily 
soluble in water; at the ordinary temperature it requires 16 parts, 
but much less quantities at the boiling-point. 

Potassium permanganate is one of the most powerful oxidizing 
agents with which we are acquainted, and attacks a number of bodies 
which are not affected by other oxidizing substances. It liberates 
chlorine from hydrochloric acid, oxidizes oxalic acid to carbonic 
acid, ferrous salts to ferric salts, sulphurous acid to sulphuric acid, 
&c. If the oxidizing solution of potassium permanganate contains 
free sulphuric acid, manganous sulphate is formed, and five atoms 
of oxygen are set free from each two molecules of the permanganate 
for the purposes of oxidation ; but if the solution is neutral, hydrated 


Detection of Manganese Compounds . 467 

manganese peroxide is produced, and only three molecules of 
oxygen are liberated from two molecules of the salt : — 

2Mn0 3 • OK + 3S0 2 (0H ) 2 = 2S0 2 • 0 2 Mn + S0 2 (0K) o + 3H.,0 + 5 O. 
2Mn0 3 -OK + H 2 0 = 2 Mn 0 2 + 2 KOH + 3 O. 

If finely powdered potassium permanganate is dropped into 
concentrated sulphuric acid, a green solution is obtained, and oxy¬ 
gen accompanied by considerable quantities of ozone is evolved. 

Owing to its powerful oxidizing properties, potassium permanga¬ 
nate cannot be filtered through paper ; the paper is at once turned 
brown, and a reduced solution of a green or brown colour runs 
through. And the facility with which potassium permanganate 
oxidizes organic substances generally makes it well adapted for 
purposes of disinfection. A solution of the crude salt or of the 
sodium compound is largely used in England under the name of 
Condfls fluid. 

The other permanganates resemble that of potassium ; they are 
all of a red colour and soluble in water. 

Detection of IVXanganese Compounds. 

The amethyst colour imparted by a small quantity of a manga¬ 
nese compound to a bead of borax in the oxidizing flame of the 
blowpipe, and the disappearance of this colour when the bead is 
brought into the reducing flame, serve to detect the metal even in 
the presence of iron. Still more delicate is the green colour due 
to a manganate produced when a minute quantity of any man¬ 
ganese compound is fused on a piece of platinum foil with sodium 
carbonate and a little potassium nitrate. The test may be con¬ 
firmed by dissolving the green mass in water, acidulating and 
warming, when the solution becomes pink (permanganate). The 
flesh-coloured precipitate of manganous sulphide produced when 
ammonium sulphide is added to a solution of a manganous salt, 
and the darkening of the white precipitate (manganous hydrate) 
thrown down by ammonia on exposure to the air, are characteristic 
reactions of manganese compounds. But the colour of the sulphide 
is obscured by a small quantity of a salt of iron, which gives black 
ferrous sulphide with the ammonium sulphide. 

Ferric salts are at once distinguished from those of manganese 
by the white precipitate of manganous ferrocyanide, produced 
when a solution of potassium ferrocyanide is added to one of a 
manganous salt. Iron and manganese, which so often occur 

h h 2 


468 Text-Book of Inorganic Chemistry. 

together, may be approximately separated by the addition of 
ammonium chloride and ammonia, which only precipitates ferric 
hydrate. The solution must be quickly filtered, otherwise man¬ 
ganic hydrate is produced by oxidation, and this is not dissolved 
by ammonium chloride. A more exact method consists in adding 
an excess of sodium acetate to the dilute solution of the two salts 
and then boiling, when the whole of the iron is precipitated as 
basic ferric acetate, while manganous acetate remains in solution. 


CHROMIUM. 

'Chemical Symbol: Cr .—Atomic Weight : 52 * 2 . 

Chromium, so-called on account of the various and beautiful 
‘colours of its compounds (xP^f JLa = colour;, is not very widely 
distributed in nature, and never occurs in the metallic state. By 
far the most important mineral containing chromium, and that 
from which its compounds are always prepared, is chrome-iron ore : 
Fe0,Cr 2 0 3 . 

Metallic chromium is only of scientific interest, and has not yet 
oeen applied to any useful purpose. It may be obtained, according 
to Wohler, by mixing one part of violet chromic chloride with two 
parts of a fused mixture of potassium and sodium chlorides, and 
heating this with two parts of granulated zinc in a Hessian crucible 
under a layer of potassium-sodium chloride. The crucible is 
heated up to the boiling-point of zinc, recognized by the bubbling 
noise, kept at this temperature for ten minutes, then removed from 
the furnace and knocked several times on the floor to cause the 
metal to unite at the bottom of the crucible. On breaking open 
the crucible when cold, a regulus of an alloy of zinc and chromium 
is obtained, which is treated with dilute nitric acid until all the 
zinc is dissolved. The chromium then remains as a pale grey, 
crystalline powder. According to Bunsen, chromium may be 
obtained by electrolyzing a concentrated solution of chromic 
chloride. It then forms brittle lustrous scales, of the colour of 
iron. The metal is not oxidized by moist air so rapidly as iron ; 
its specific gravity is about 6 - 5 . When heated in the air it becomes 
covered with a thin film of the oxide, imparting to it various colours, 
like steel; it is only slowly oxidized in dry air, even at a red-heat. 
Hydrochloric acid dissolves the metal with evolution of hydrogen ; 



Chromium. 469 

sulphuric acid acts similarly on warming. Nitric acid, even when 
hot and concentrated, does not act on it. 

Chromium, like manganese, forms several compounds with 
oxygen. We are acquainted with —chromous oxide : CrO, chromic 
oxide : Cr 2 0 3 , chromic anhydride : Cr0 3 , and perchromic acid : 
HCr0 4 (?). The first and last of these compounds are unstable, 
and scarcely known. Chromic oxide, like alumina and other 
sesquioxides, is a weak base, while chromic anhydride possesses 
strong acid properties. 


CHROMIC OXIDE AND CHROMIC SALTS. 


Chromic Oxide : Cr 2 0 3 .—This compound, which is isomor- 
phous with alumina and ferric oxide, is amorphous or crystalline 
according to its mode of preparation, but always possesses a green 
colour. Chromic oxide may be obtained by a variety of methods, 
of which the following are the most important:— 

(i.) By heating chromic hydrate. 

(ii.) By heating mercurous dichromate to redness :— 


O 


| Cr0 2 -OHg 
I CrOo - OHg 


= Cr 2 0 3 + 2 Hg + 20 2 


when it remains as a dark green powder. 

(iii.) By igniting ammonium dichromate :— 


O 


Q&O nh; = Cr ‘°3 + < H *o + N *> 


when the chromic oxide remains behind in olive green scales. 

(iv.) By very strongly heating a mixture of common salt and 
potassium dichromate covered with a layer of common salt in a 
Hessian crucible. On extracting the fused mass with water, the 
chromic oxide remains as glistening iridescent crystals 


O 


Cr0 2 *0K 

Cr0 2 -0K 


Cr 3 0 3 + K 2 0 + $0, 


the common salt only serving to assist the crystallization of the 
chromic oxide. 

(v.) Finally, by passing the vapour of chromyl chloride through 
a porcelain tube heated to redness :— 

2 Cr0 2 Cl 2 - Cr 2 0 3 + 2 CI, + 0, 



470 Text-Book of Inorganic Chemistry. 

when chlorine and oxygen are given off, and the chromic oxide is 
deposited on the walls of the tube in lustrous dark-green crystals, 
which are so dark in colour as to appear almost black. These 
crystals are as hard as corundum, and have a specific gravity of 5 ’ 2 . 

Strongly ignited chromic oxide is insoluble in acids, even in 
concentrated sulphuric acid. It may be heated either in the air 
or in hydrogen to the highest temperatures of our furnaces without 
change; but is reduced when strongly heated with charcoal. 
When heated with potassium (or sodium) carbonate and exposed 
to the air, or with the addition of a little nitre, chromic oxide gives 
soluble potassium (or sodium) chromate. Insoluble compounds 
containing chromium— eg. chrome-iron ore—are thus brought into 
solution. 

Chromic oxide easily dissolves in molten glass as well as in 
borax, and the chromic silicate or borate so produced possesses a 
beautiful green colour. On account of this property and because 
it remains unchanged at a high temperature, chromic oxide is 
largely used for colouring glass green and for painting on porce¬ 
lain. 

Chromic oxide, like alumina and other sesquioxides, is a weak 
base. It unites with acids and with bases, forming, in the latter 
case, compounds which correspond to the aluminates. These 
compounds, sometimes called chromites , include chrome-iron ore : 

Fe0,Cr 2 0 3 = J 0 2 Fe", referred to above, and which is distin¬ 
guished by the difficulty with which it is attacked by most acids. 
Among other chromites are the soluble green compounds pro¬ 
duced by digesting freshly precipitated chromic hydrate with 
caustic potash or soda. These compounds are, however, less 
stable than the corresponding aluminates, and are decomposed on 
boiling their solutions. It is remarkable that chromic hydrate 
when mixed with ferric hydrate is not dissolved by caustic al¬ 
kalies. 

Chromic Hydrate : Cr(OH) 3 .—This substance separates when 
a solution of a chromic salt is mixed with ammonia as a bluish-green 
bulky precipitate, which when dried forms a dirty-green powder. If 
dried at ioo° it possesses the above composition. When gradually 
heated more strongly it begins to glow throughout its entire mass 
at one particular temperature, and is then converted into dark-green 
chromic oxide, which is now insoluble in acids. Freshly precipitated 


47i 


Chromic Sulphate . 

chromic hydrate easily dissolves in acids, and produces two kinds 
of differently coloured chromic salts ; the solutions of the one kind 
are green, and those of the other violet. The violet salts change 
into green by continued boiling, but on allowing the latter to stand 
for some time a change takes place from the green to the violet. 
The violet salts crystallize easily ; but the green solutions, when 
evaporated down, leave a green amorphous mass . 1 Chromic 
hydrate, precipitated by ammonia from a solution of a violet salt, 
is somewhat soluble in an excess of the reagent, but is rendered 
completely insoluble on boiling. 

A hydrate of chromium of a fine green colour is known in 
trade as GuignePsgreen, and has the composition : Cr 2 0 3 , 2H 2 0 = 

O \ It is obtained by fusing a mixture of potassium 

(Cr((Jri ; 2 

dichromate and crystallized boric acid, and then extracting the 
fused mass with water. During the reaction, potassium and 
chromic borates are produced, and the latter compound is decom¬ 
posed when treated with water. 

Chromic Sulphate: (S0 2 ) a 0 6 Cr 2 + I 5 H 2 0 .—This salt crys¬ 
tallizes in violet octahedra, when chromic hydrate is dissolved in 
sulphuric acid and allowed to stand for some time. If its solution 
is mixed with potassium sulphate, the two sulphates unite to form 

chrome-alum : OK f + I2H 2°> which crystallizes out in fine 

dark-purple octahedra when the solution is allowed to evaporate 
slowly. 

Chrome-alum is, however, best prepared by the reduction of a 
mixture of potassium dichromate and sulphuric acid by some easily 
oxidized substance, such as sulphurous acid or alcohol. The 
chromic acid which is thus set free by the sulphuric acid parts 
with oxygen, and is reduced to chromic oxide. One of the best 
reducing substances to employ is ordinary alcohol, which then 
becomes oxidized to the volatile compounds—aldehyde and acetic 
acid. To obtain chrome-alum in this way, 2 parts of a saturated 
solution of potassium dichromate are mixed with 3 parts of con¬ 
centrated sulphuric acid in a large flask, and alcohol added drop 
by drop from a separating funnel, with continued agitation, until 

1 The chemical difference between the violet and green modifications is at. 
present unknown. It has sometimes been thought that the green varieties 
contain a mixture of an acid and basic salt. Ed. 


472 


Text-Book of Inorganic Chemistry. 


the red liquid has changed to a dark green. On allowing the 
green liquid to stand for some time in an open dish it gradually 
becomes of a violet colour, and, in consequence of evaporation, 
small octahedra of chrome-alum begin to be deposited. If these 
small crystals are allowed to lie in the solution, which is kept in a 
cellar where the temperature is nearly uniform, and are turned 
daily, they gradually grow to large regular octahedra. The above 
reaction takes place according to the following equation : 



(Cr0 2 -OK 

(Cr0 2 -OK 


in which the oxidation of the alcohol by the three atoms of oxygen 
is not shown. 

As in ordinary alum, the potassium in chrome-alum may be 
displaced by any of the alkali-metals, or by thallium, without 
changing its crystalline form. 

Chromic Nitrate : (N0 2 ) 3 0 3 Cr + qH.O, crystallizes with diffi¬ 
culty from the violet solution obtained by dissolving chromic 
hydrate in nitric acid. 

Chromic Phosphate : PO • 0 3 Cr + 6H 2 0, falls as a dirty-green 
precipitate when a solution of sodium phosphate is added to one 
of chrome-alum. It becomes crystalline on standing under the 
liquid. 

Chromic chloride : Cr 2 Cl 6 .—This chloride sublimes in beau¬ 
tiful peach-blossom coloured scales, when a mixture of chromic 
oxide and charcoal is heated in a stream of dry chlorine ; it is 
distinguished from the corresponding chlorides of aluminium and 
iron by its insolubility in water. When heated in the air it loses 
chlorine, and takes up oxygen, being first converted into an oxy¬ 
chloride and then into chromic oxide. 

A simple method of preparing considerable quantities is as 
follows. The dried mixture of chromic oxide, charcoal, and starch 
paste is placed in a Hessian crucible, on which is cemented a 
second crucible of similar size placed upside down, and through 
the bottom of which a hole is bored to receive the glass tube 
bringing the chlorine. The two crucibles are placed in a char¬ 
coal furnace so that the lower one is surrounded by the glow¬ 
ing charcoal, and is heated to redness, while the upper one is 
only moderately heated. Chlorine is then led into the upper 
crucible through a glass tube which fits the opening loosely, and 


Chromic Chloride. 


473 


the lower crucible is raised to a bright red-heat. The carbonic 
oxide and the excess of chlorine escape through the opening at 
the top, and the chromic chloride condenses on the sides of the 
upper and cooler crucible : a portion of the chloride remains in 
the lower crucible loosely attached to the carbonaceous residue, 
from which it may be easily removed when cold. 

Chromic chloride when pure is quite insoluble in water, but 
dissolves at once to form a green liquid when the water contains a 
trace of crystalline chromous chloride. Mere contact with the 
chromous chloride suffices to convert the insoluble violet chromic 
chloride into the soluble green variety. 

By dissolving chromic hydrate in hydrochloric acid and eva¬ 
porating down, a green hydrated chromic chloride is obtained, 
which loses water and hydrochloric acid when heated, and leaves 
a basic chromic chloride behind. 

Chromous Chloride : CrCl 2 .—If chromic chloride is heated to 
low redness in a stream of hydrogen it loses chlorine, and is con¬ 
verted into a white crystalline mass of chromous chloride. This 
chloride dissolves in water to a blue liquid, which then absorbs 
oxygen when exposed to the air, and produces an oxychloride. 
Chromous chloride, like ferrous chloride, can dissolve considerable 
quantities of nitric oxide, forming a dark-brown solution. The 
merest trace of this chloride can convert considerable quantities of 
the violet chromic chloride into the soluble green modification (see 

When caustic potash is added to a solution of chromous 
chloride a brown precipitate is produced, which is probably of 
chromous hydrate : Cr(OH) 2 , but which possesses so strong an 
attraction for oxygen, that it decomposes water and forms tn- 
chromic tetroxide : Cr 3 0 4 , with evolution of hydrogen. 


CHROMIC ACID AND THE CHROMATES. 

Chromic Anhydride Cr0 3 . 

True chromic acid is unknown in the free state; when 
separated from its salts it breaks up into water and chromic 
anhydride. 




474 Text-Book of Inorganic Chemistry. 

Chromic anhydride, commonly known as chromic acid, is 
easily obtained by adding ij volume of concentrated sulphuric 
acid to i volume of a saturated solution of potassium dichromate 
at 5 o c . The mixture becomes strongly heated, and deposits on 
cooling long bright red prisms of the chromic acid liberated by 
the sulphuric acid. When quite cold the mother-liquor, containing 
acid potassium sulphate, is poured off, and the crystals dried first on 
an asbestos filter, and then by placing them on a porous plate of 
unglazed porcelain, which absorbs the remaining traces of water. 
By recrystallizing from a small quantity of water and repeating the 
drying process, the crystals may be obtained pure and free from 
sulphuric acid. 

Chromic acid is easily soluble in water and deliquesces in moist 
air. The solution is of a reddish-brown or yellow colour according 
to its strength, and reacts strongly acid. Chromic acid contains 
one half of its oxygen very loosely combined, and is decomposed 
into chromic oxide and oxygen not only when heated alone : 
2Cr0 3 = Cr 2 0 3 + 3 O, but also when heated with sulphuric acid, 
chromic sulphate being formed in this case. In the same way it 
evolves chlorine with strong hydrochloric acid :— 

2 CrO a + 12 HCI = Cr 2 Cl 6 + 6H 2 0 + 3 d , 1 

It oxidizes organic substances so powerfully that few can with¬ 
stand its action. If the solution is filtered through paper, or if the 
crystals are dried with paper, the chromic acid at once becomes of 
a green colour from chromic oxide. Alcohol when dropped on the 
crystalline substance is so rapidly oxidized that it catches fire. 

If a solution of chromic acid is mixed with hydrogen peroxide 
a compound of a beautiful blue colour is formed, which easily 
dissolves in ether. It has been suggested that this compound is 
fierchromic acid, corresponding to permanganic acid, but its com¬ 
position is doubtful (p. 91 ). 

The Chromates. 

[ OH 

The hypothetical chromic acid: Cr0 2 j q is, like sulphuric 
acid, a dibasic acid, but is distinguished from this and all other 

1 An excellent method of obtaining a long-continued, regular stream of 
chlorine consists in gently heating potassium dichromate with an excess of 
hydrochloric acid.—E d. 


Chromic Acid and the Chromates . 


475 


dibasic acids by the fact that it never produces acid salts. That 
acid salts of chromic acid might be produced cannot be doubted, 
but we are as yet unacquainted with the conditions under which 
they can exist. They decompose as soon as produced into the 


normal salts of dichromic acid 


0 (Cr0 2 *0H an< ^ water * If) f° r 


I OK 

example, normal potassium chromate : Cr0 2 jQjr is mixed with a 

solution of chromic acid, or with a stronger acid, then it is probable 
that acid potassium chromate is first produced :— 


or 


Cr0 3 


(OK 

(OK 


+ 


Cr0 2 


OH 
OH * 


2Cr0 2 


(OH 

(OK 


CrO 


JOK 

2 {ok 


+ 


so. 


(OH 

(OH 


~ ~ (OH 
Cr ° 2 (OK 


so 2 


(OH 

{ok 


and that two molecules of the acid chromate immediately afterwards 
break up into the normal dichromate and water :— 


2 CrO. 


(OH 

{OK 


n (Cr0 2 -0K 
U {Cr0 2 -OK 


h 2 o. 


These soluble dichromates are again converted into the normal 
chromates on the addition of a base :— 


°{crOl'ol + 2KOH = 2Cr °4o£ + H ’°‘ 

Similar relations to that between chromic acid and dichromic 
acid also exist between sulphuric and disulphuric acids (p. 167 ), and 
between phosphoric and pyrophosphoric acids (p. 219 ) ; the only 
difference being that the normal salts of disulphuric and pyrophos¬ 
phoric acids are produced from the acid sulphates and acid phos¬ 
phates respectively at a high temperature, while the change from 
the acid chromates to the normal dichromates takes place at the 
ordinary temperature. It is possible that acid chromates may exist 
at a low temperature. 


Potassium Dichromate : 


K 2 Cr 2 0 7 


(Cr0 2 *0K 
{Cr0 2 -OK 


—This 


salt 


is manufactured on a large scale from chrome-iron ore : FeO, 
Cr 2 0 3 , and is the starting-point for the preparation of all other 
chromium compounds. It is obtained when the finely powdered 
ore is mixed with potassium carbonate and nitre and fused for some 
time. The iron is then oxidized to ferric oxide and the chromium 


476 Text-Book of Inorganic Chemistry . 

to chromic acid, which unites with the potassium and forms yellow 
potassium chromate. The latter compound goes into solution on 
digesting the fused mass with water. On the large scale the mixture 
of the pulverized ore and potassium carbonate, without nitre, is 
heated in a reverberatory furnace and the oxidation effected by 
excess of air. A certain quantity of lime is usually added, which 
keeps the mass porous and prevents complete fusion. The yellow 
solution of potassium chromate obtained by afterwards extracting 
the mass with water is then acidulated with sulphuric acid, which 
both destroys any excess of potassium carbonate and converts the 
normal chromate into red potassium dichromate, and the solution 
is allowed to crystallize. The chromate is converted into the dichro¬ 
mate because the latter salt crystallizes much more easily from its 
solutions, and can, therefore, be more readily purified. 

Potassium -dichromate, purified by recrystallization, forms 
orange-red tabular crystals of the above composition, and without 
water of crystallization. It dissolves in ten parts of water at the 
ordinary temperature, but is more soluble in hot water ; the solu¬ 
tion reacts faintly acid. When heated to low redness it melts 
unchanged to a dark red liquid, and at very high temperatures is 
decomposed. Concentrated sulphuric acid separates crystalline 
chromic acid from its aqueous solution. Sulphuretted hydrogen 
reduces a solution of potassium dichromate and produces a preci¬ 
pitate of chromic hydrate and sulphur, while the potassium is 
converted into sulphide or sulphydrate, which remains in solu¬ 
tion :— 

° jcrCVOK + 4H 2 S - 2 Cr(OH) 3 + 3 S + K 2 S + H 2 0. 

No precipitate of chromic hydrate is produced if hydrochloric 
or sulphuric acid is previously mixed with the solution of potassium 
dichromate ; the red colour of the solution then changes to a green 
from the chromic chloride or sulphate produced. 

If a solution of potassium dichromate is mixed with caustic 
potash or potassium carbonate until the red colour has become a 
pure yellow, and the solution evaporated down, 

Potassium chromate : Cr0 2 (0K) 2 , separates out in yellow 
crystals, isomorphous with potassium sulphate. The salt is very 
easily soluble in water, and cannot therefore be so easily crys¬ 
tallized as the dichromate. Its yellow solution reacts slightly 
alkaline, and becomes red on the addition of an acid owing to the 


Chromic Acid and the Chromates. 4 77 

formation of potassium dichromate, which in its other properties 
it closely resembles. 

Sodium Chromate. —Both sodium chromate and dichromate 
closely resemble the corresponding potassium salts, except that 
they are much more soluble in water, and even deliquesce in the 
air. 

Ammonium Dichromate : O q.q ' q^H 4 cr 7 sta ^i zes i n yellow 

scales, and otherwise resembles the potassium salt. It is easily 
soluble in water, and gives green chromic oxide when heated 

(p. 469 ). 

Ammonium Chromate'. Cr0 2 (0NH 4 ) 2 , crystallizes in yellow 
needles on evaporating a mixture of aqueous chromic acid and 
excess of ammonia. The crystals remain unchanged in the air. 

Calcium and Strontium Chromates crystallize in yellow prisms, 
and are much less soluble in water than the potassium salt. 
Barium Chromate : Cr0 2 -0 2 Ba, is nearly insoluble in water, and 
separates as a pale-yellow precipitate when a solution of potassium 
chromate pr dichromate is added to one of bariunTchloride. It is 
used as a yellow pigment. 

The chromates of the heavy metals are coloured substances, 
insoluble in water. 

Chlorchromic Acid and Cbromyl Chloride. 

Chromic acid, which resembles sulphuric acid in so many 

( Cl 

points, also forms a compound, chromyl chloride : CrO j C1 corre- 

f rq 

sponding to sulphuryl chloride : S0 2 1 C1 in which the two atoms 

of hydroxyl of the hypothetical acid are displaced by chlorine. A 

f Cl 

second compound, corresponding to chlorsulphonic acid: S0 2 j qj_j 

and in which only one atom of hydroxyl is displaced, is known in 
the salts of chlorchromic acid— e.g. potassium chlorchromate : 

CrO - The acid itself has not yet been prepared. To obtain 

(OK 

this potassium salt, 3 parts of finely powdered potassium dichro¬ 
mate are mixed with water to a paste, then gently warmed with 
4 parts of concentrated hydrochloric acid until a faint odour of 


478 


Text-Book of Inorganic Chemistry. 


chlorine is perceptible. The salt afterwards separates in large 
red, tabular crystals as the clear, red solution cools 

°{crO:'ol + 2HC1 " 2Cr °4oK + 

The salt is decomposed by water, but dissolves in hydrochloric 
acid unchanged. When gently heated it evolves chlorine copiously. 


Chromyl Chloride: Cr 0 2 j^| is a blood-red, heavy liquid of 

i'9 specific gravity, and boiling at ii6 °. It possesses a powerful 
piercing odour, and fumes in the air. In order to prepare this 
compound, io parts of common salt are fused together with 12 
parts of potassium dichromate, the fused mass poured on to a 
clean iron plate and broken into pieces when cold. It is then in¬ 
troduced into a large retort connected with a condenser, and 40 
parts of concentrated sulphuric acid, containing some Nordhausen 
acid, added. A brisk reaction ensues, which must afterwards be 
supported by a gentle heat, and the compound distils over into the 
well-cooled receiver. It is purified by repeated distillation, finally 
in a stream of carbonic acid gas. The reaction is represented by 
the following equation :— 


O 


CrOo’OK 
CrCVOK 

2 SO 


+ 4NaCl + 6 S 0 2 j 


OH 

OH 


(OK 
I OH 


.co l° Na 

4S °2]OH 


= 2 Cr 0 2 CL 

3H 2 0. 


+ 


According to which 1 molecule of potassium dichromate and 
4 molecules of sodium chloride give, when heated with an excess 
of sulphuric acid (about 9 molecules), 2 molecules of chromyl 
chloride, 2 of acid potassium sulphate, 4 of acid sodium sulphate, 
and 3 of water, for the retention of which at least 3 molecules of 
sulphuric acid are required. 

Chromyl chloride is decomposed by water into chromic and 
hydrochloric acids. It acts energetically on phosphorus, sulphur, 
and many organic substances with the production of light and heat. 
When its vapour is led through a red-hot tube it is decomposed into 
chlorine, oxygen, and chromic oxide. If the requisite quantity of 
chromyl chloride is dropped into a solution of normal potassium 
chromate, containing a little acetic acid, the liquid becomes filled 
with crystals of potassium chlorchromate :— 


Cr 0 2 Cl 2 4 CrO a (OK) 2 = 2CrO a j q K 


Detection of Chromium Compounds. 


479 


Detection of Chromium Compounds. 

Compounds of chromium, in any form, may be readily and cer¬ 
tainly detected by simple blowpipe tests. When heated with 
sodium carbonate and a little nitre on a piece of platinum foil, they 
all produce yellow sodium chromate, and the fused mass, when dis¬ 
solved in water and acidulated with acetic acid, gives a yellow pre¬ 
cipitate of lead chromate on the addition of a few dtops of a 
solution of lead acetate. 

A trace of a chromium compound imparts a green colour to a 
borax bead (due to chromium borate), both in the oxidizing and 
reducing flame, but which becomes converted into yellow sodium 
chromate when heated for a long time in the oxidizing flame. 

An aqueous solution of a chromate is recognized by the yellow 
precipitate of lead chromate produced on the addition of a solution 
of lead acetate ; and by the change of the yellow or red solution to 
a green salt of chromic oxide when acted upon by sulphuretted 
hydrogen, sulphurous acid, alcohol, or other reducing agents in the 
presence of hydrochloric or sulphuric acid. 

The soluble salts of chromic oxide all yield chromic hydrate 
with ammonium sulphide, ammonia, or caustic soda, and are not 
precipitated by sulphuretted hydrogen. In this respect the salts 
of chromic oxide resemble those of alumina. Freshly precipitated 
chromium hydrate, like aluminium hydrate, is easily dissolved by 
caustic soda, but the hydrate is reprecipitated on boiling, or when 
allowed to stand. 

Compounds of chromic oxide may therefore be separated from 
those of alumina by adding excess of caustic soda and boiling; 
they may be separated from salts of iron by fusing with sodium 
carbonate and nitre, when the iron is converted into insoluble ferric 
oxide, while the chromium takes the form of a soluble chromate, 
which may be extracted by digesting the fused mass with water. 


URANIUM. 

Chemical Symbol : U .—Atomic Weight'. 240 . 

This somewhat rare element is never found free in nature; it 
occurs almost exclusively as the oxide : U 3 0 8 = U0 2 ,2U0 3 , in the 
mineral pitchble7ide. The metal may be obtained by fusing 



480 Text-Book of Inorganic Chemistry. 

uranous chloride with sodium, best in an iron crucible, and then 
forms either a black powder or a hard metallic button of a grey 
colour. It melts at a red heat, and has a specific gravity of 187 . 
The metal does not decompose water even at ioo°, but oxidizes 
when heated in the air, and is easily dissolved by dilute acids. 

Uranium unites in two proportions with oxygen, forming ura¬ 
nous oxide : U 0 2 , and uranic oxide : U 0 3 , and a third oxide also 
exists of the composition : U 3 O g , which may be considered as a 
compound of the first two: ’U0 2 + 2U0 3 . Pitchblende consists 
chiefly of this oxide, but contains also compounds of lead, copper, 
arsenic, iron, manganese, cobalt, and other metals. To obtain 
pure uranium compounds from this mixture the following process 
may be adopted. The powdered mineral is heated with mode¬ 
rately concentrated sulphuric acid and small quantities of nitric 
acid added from time to time. When only a white insoluble 
residue (of lead sulphate, &c.) remains, the liquid is evaporated 
down to expel most of the acid, then diluted with water and the 
clear solution saturated with sulphuretted hydrogen to precipitate 
lead, copper, arsenic, &c. The solution is next filtered off from 
these sulphides, the iron oxidized again by nitric acid, and 
ammonia added in excess. The dark-yellow precipitate so 
obtained, which chiefly consists of ferric hydrate and ammonium 
uranate, is then well washed and treated with a warm concentrated 
solution of ammonium carbonate containing free ammonia, until 
all the uranium is dissolved and the precipitate consists entirely of 
dark-brown ferric hydrate. The yellow solution, which must be 
quickly filtered, deposits small, hard, yellow crystals of ammonium- 
uranic carbonate, and a further deposit of the same salt is obtained 
as a yellow powder when the solution is boiled. This compound 
when heated to redness leaves the pure dark-green oxide : U 3 O s . 

Uranic Nitrate'. ^q~ j 0 2 (U0 3 ) + 6H 2 0, is obtained in large 

yellow deliquescent prisms when the preceding or any other oxide 
of uranium is dissolved in nitric acid and the solution evaporated 
down. If the salt is heated to 250 ° until nitrous fumes are no 
longer evolved, it becomes changed into 

Uranic Oxide : U0 3 = (U0 2 )0, which remains behind as a 
yellow powder, and which when strongly heated loses oxygen and 
reproduces uranoso-uranic oxide : U 3 0 8 . 

Uranic Sulphate'. S0 2 -0 2 (U0 2 ) + 3 H 2 Q, may be obtained by 


Uranium. 481 

decomposing the nitrate with sulphuric acid, and crystallizes in 
pale yellow needles. 

Uranic Carboyiate. —Ammonium carbonate, when added to a 
solution of a uranium compound, produces a yellow precipitate, 
which, however, is not uranium carbonate, but a double carbonate 
of ammonium. This salt is soluble in an excess of ammonium 
carbonate, and if this solution is allowed to evaporate in the air, or 
if gently warmed, it deposits small yellow glistening crystals of the- 
compound: CO *0 2 (U0 2 ) + 2 CO • (ONH 4 ) 2 , which is only slightly 
soluble in water. This double salt is used to separate uranium 
from other metals. 

Uranic Phosphate : PO j + 4H 2 0, is thrown down as a 

yellow, crystalline precipitate on adding a solution of phosphoric 
acid or a soluble phosphate to one of uranic acetate. It is in¬ 
soluble in acetic acid. If the solution of uranium contains ammo¬ 
nium chloride, the precipitate which is produced on the addition 

of a phosphate has the composition: PO j This com- 

(OJN H 4 

pound is also insoluble in acetic acid, and is used to estimate phos¬ 
phoric acid, especially in manures. 

The mineral auticnite is a double phosphate of uranium and 
calcium. 

From the composition of the uranium compounds given above, 
it is clear that uranic oxide is not a triacid base like alumina, ferric 
oxide, and chromic oxide, but that it always combines only with two 
molecules of a monobasic acid, like lime, magnesia, and the oxides 
of other dyad metals. Two atoms of oxygen are firmly united 
with an atom of uranium to form a dyad radical : U0 2 , called 
uranyl , and the salts of the oxide— e.g. uranic nitrate, are properly 
called uranyl salts— e.g. uranyl nitrate. This rule, however, is not 
followed, because no compound of the composition : (N0 2 ) 6 0 6 U, is 
known. If, then, uranic oxide : U0 3 , is to be considered as uranyl 
oxide : (U0 2 )0, we should expect that on dissolving this oxide in 
hydrochloric acid a chloride of the composition : (U0 2 )C1 2 , would 
be obtained, According to the equation :— 

(U0 2 )0 + 2 HCI = (U0 2 )C1 2 + h 2 o, 

and as a matter of fact the compound : Ura?iyl chloride , which 


482 Text-Book of Inorganic Chemistry . 

crystallizes with difficulty from the yellow solution obtained by 
dissolving uranic oxide in hydrochloric acid, has this composition. 

Uranic oxide unites with strong bases to form a series of com¬ 
pounds called uranates , in which it plays the part of an acid. All 
these salts, even those of the alkalies, are insoluble in water. 

Potassium uranate'. 0 [yC^-OK corres P ondin § to potassium 

dichromate, is obtained as a yellow precipitate on adding caustic 
potash to a solution of uranium nitrate. It bears a high tempera¬ 
ture without decomposition. Sodium uranate and Ammonium 
uranate closely resemble the potassium salt, and are obtained by 
adding caustic soda or ammonia to a solution of uranium nitrate. 
Ammonium uranate dissolves in ammonium carbonate in the cold, 
and when ignited in the air leaves uranoso-uranic oxide as a green 
powder. 

Uranyl Sulphide : (U0 2 )S, is precipitated as a black amor¬ 
phous powder, nearly insoluble in water, when ammonium sulphide 
is added to a solution of an uranic salt. It is soluble in ammonium 
carbonate, and is therefore not thrown down if the solution to 
which the ammonium sulphide is added contains this salt. 

Uranous Oxide : U0 2 , which is produced by glowing uranic 
oxide or uranoso-uranic oxide in a stream of hydrogen, is a black 
or dark brown crystalline powder, insoluble in dilute hydrochloric 
acid, but dissolving in nitric acid and in concentrated sulphuric 
acid. If a solution of uranous oxide in sulphuric acid, which is 
of a green colour and contains uranous sulphate : (S0 2 ) 2 0 4 U, is 
mixed with caustic soda, a brown precipitate of uranous hydrate : 
U(OH) 4 , is produced, which turns black on boiling. 

Uranous Chloride : U Cl 4 , may be prepared by igniting a 
mixture of any oxide of uranium and charcoal in a stream of chlorine, 
and then sublimes in lustrous dark green octahedra. It is readily 
soluble in water and deliquesces in moist air. 


Uranoso-uranic Oxide: U 3 0 8 = U0 2 ,2U0 3 , is obtained as a 
dark green powder, by glowing either uranous or uranic oxide in 
the air. In the former case oxygen is absorbed, in the latter it is 
given off. It is easily soluble in nitric acid, but only dissolves in 


Cobalt and Nickel. 


483 

sulphuric or hydrochloric acid with difficulty, and withstands a high 
temperature without change. 

Glass (or a borax bead) is coloured a bright green by this oxide, 
while uranic oxide imparts a greenish yellow tint, which fluoresces 
strongly in bright light. 


COBALT AND NICKEL. 

These twin metals, as they may be called, always occur in 
association with one another in nature, the one being in excess in 
some minerals, and the other in others. Both are constant 
constituents of metallic meteorites, though the nickel is always in 
excess. The two metals have nearly the same atomic weight, and 
their compounds are so much alike, chemically, that for a long 
time no method was known of separating them. Their salts are 
chiefly distinguished from one another by their colour : those of 
cobalt are red or blue, and those of nickel green. 

The similarity between the two metals extends even to their 
names. The external appearance of the cobalt ores led to the 
belief that a valuable metal might be extracted from them by the 
ordinary metallurgical processes. But when all attempts in this 
direction proved futile, the German miners imagined the ores to be 
possessed of an evil spirit or Kobold (= sprite, goblin), and this 
word then gave the name cobalt to the metal afterwards discovered 
in the ores. In the same way, many attempts were made to ex¬ 
tract copper from the commonest ore of nickel—Kupfer-nickel— 
which is itself of a copper colour ; but when all these endeavours 
were in vain the ore was thrown on one side and called Nickel (an 
abusive word for an obstinate person in some South German 
dialects), a name which has now been retained for the metal itself. 

Of the two metals, cobalt was first used in the arts, and in the 
form of its oxide, after it had been discovered that this compound 
dissolves in molten glass with a dark-blue colour, and that such 
glass when finely powdered yields a fine blue pigment {smalt). In 
the manufacture of smalt the nickel compounds are separated 
under the name of cobalt-speiss, and this was formerly thrown 
away as useless. Atone time this speiss was used in Hesse for 
mending the roads, but when its value became known the road 
scrapings were carefully collected and worked for nickel. 


1 1 2 



484 Text-Book of Inorganic Chemistry. 

Cobalt is only used technically in the form of its oxide, not as 
the metal ; while nickel, on the other hand, is almost exclusively 
used in the metallic state for the preparation of alloys and for 
nickel-plating. 


COBALT. 


Chemical Symbol : Co .—Atomic Weight : 59- 

Cobalt is only found native in metallic meteorites, and then in 
very small quantities. The most important minerals which contain 
the metal are smaltite or speiss-cobalt : CoAs 2 , and glance-cobalt : 
CoS 2 ,CoAs 2 = CoAsS. From these ores, which usually contain 
copper, iron, and nickel, cobalt or a pure cobalt compound may 
be obtained by a variety of methods. To prepare cobalt from 
either of the above ores on the small scale, the following pro¬ 
cess may be adopted. The ore is first freed from some of its 
sulphur and arsenic by roasting in the air, then dissolved in 
hydrochloric acid, with the addition of small quantities of nitric 
acid, until completely decomposed. Sulphuretted hydrogen is 
next passed through the acid liquid to precipitate arsenic, copper, 
lead, &c., and after the iron has been again oxidized by nitric 
acid, a solution of sodium carbonate added until just neutral, 
by which all the iron is precipitated as ferric hydrate. The 
filtered solution, on further addition of sodium carbonate, gives 
a precipitate of the mixed carbonates of cobalt and nickel, which, 
when washed and digested with a strong solution of oxalic acid, 
are converted into the insoluble oxalates, both of which dissolve 
in strong ammonia. This solution of the oxalates in ammonia 
is then allowed to stand in the air for some time, so that the 
ammonia may slowly evaporate, when the whole of the nickel is 
precipitated as oxalate, while the cobalt remains in solution. The 
clear purple solution, which now only contains cobalt, is evaporated 
to dryness, the residue glowed, and then dissolved in hydrochloric 
acid, and after the excess of hydrochloric acid has been removed 
by heating, precipitated with a solution of oxalic acid. The cobalt 


f c o 

oxalate : j 0 2 Co, so obtained is decomposed when heated into 

carbonic acid and cobalt, the latter fusing at a sufficiently high 
temperature. To reduce the metal from the oxalate, a narrow, 



485 


Compounds of Cobalt. 

deep, unglazed, porcelain crucible is filled with the dry salt, then 
covered, placed in a larger Hessian crucible lined with charcoal, 
and exposed to a high temperature in a wind furnace. The metal 
then melts to a steel-grey button, with a tinge of red. It has a 
specific gravity of 8'6, is malleable, and can be highly polished. It 
is only fused with difficulty, is more compact than iron, and is 
slightly magnetic. Cobalt is less readily oxidized than iron, but 
when heated in the air becomes covered with a film of the oxide ; it 
dissolves in hydrochloric acid and dilute sulphuric acid, but less 
readily than iron, and then produces cobaltous chloride and sul¬ 
phate respectively ; nitric acid also dissolves it, forming cobaltous 
nitrate. 

Cobalt unites in two proportions with oxygen, forming cobaltous 
oxide : CoO, and cobaltic oxide : Co 2 0 3 . The former is a strong 
base, while the latter possesses more the character of a peroxide. 

Cobaltous Oxide: CoO; Cobaltous Hydrate : Co(OH) 2 . 

When caustic soda is added to a solution of cobaltous chloride, 
a avender coloured precipitate of a basic chloride is formed , the 
precipitate changes to pink cobaltous hydrate when boiled, but 
soon absorbs oxygen from the air, and becomes of a brownish 
colour. When this precipitate is heated out of contact with the air, 
it loses water and is converted into cobaltous oxide, forming a dull 
green powder. The same compound is produced when cobaltous 
carbonate is heated out of contact with air. When heated in the 
air, cobaltous oxide is converted into the three-quarter oxide—tri- 
cobaltic tetroxide : Co 3 0 4 , corresponding in composition to the 
magnetic oxide of iron. Cobaltous oxide dissolves in hydrochloric, 
sulphuric, or nitric acid, and produces bright red solutions. 

Cobaltic Oxide : Co 2 0 3 ; Cobaltic Hydrate : Co(OH) 3 . 

Cobaltic oxide is a black powder, insoluble in water, and is 
obtained by carefully heating cobaltous nitrate or cobaltic hydrate. 
The hydrate is produced when chlorine is led into an alkaline solu¬ 
tion containing cobaltous hydrate in suspension. Such a solution 
is easily obtained by precipitating a cobaltous salt with excess of 
caustic soda. The reaction is as follows 

Co(OH ) 2 + NaOH + Cl = Co(OH) 3 + NaCl. 

Cobaltic hydrate may also be obtained by precipitating a solu¬ 
tion of a cobaltous salt with sodium hypochlorite, or a solution of 
bleaching powder. Like the oxide, it is also of a black colour, and 


486 Text-Book of Inorganic Chemistry. 

both compounds dissolve easily in hydrochloric acid. This solution 
appears at first to contain cobaltic chloride : CoCl 3 , as but little 
chlorine is set free in the cold ; if, however, the liquid is gently 
warmed, chlorine is liberated in abundance, and the solution now 
contains red cobaltous chloride. 

Cobaltic oxide, when strongly heated in the air, loses oxygen 
and is converted into the oxide, Co 3 0 4 . 

Cobaltous Sulphate : S0 2 -0 2 Co + 7H 2 0, is easily soluble in 
water, and separates from its aqueous solution in dark-red crystals, 
isomorphous with ferrous sulphate. 

Cobaltous Nitrate : (N0 2 ) 2 0 2 Co + 6H 2 0, crystallizes with diffi¬ 
culty in dark-red prisms, which deliquesce in the air. When heated 
it easily parts with oxygen and all its nitrogen as nitric peroxide, 
and leaves a residue of cobaltic oxide, which becomes cobaltoso- 
cobaltic oxide : Co 3 0 4 , when more strongly heated. 

Cobaltous Nitrite is an unstable compound, and has not been 
prepared pure. If a neutral aqueous solution of a cobaltous salt 
is mixed with one of potassium nitrite in excess, a yellow, crystal¬ 
line precipitate soon separates, which is a double nitrite of potas¬ 
sium and dyad cobalt, and has the composition :— 

(NO) 2 0 2 Co' v + 2 NO-OK + H 2 0. 

This compound is slightly soluble in cold water, and dissolves 
more readily when warm, forming a red solution. 

A second double nitrite, but containing triad cobalt, is obtained 
when potassium nitrite is added in excess to a solution of a cobal¬ 
tous salt strongly acidified with acetic acid. A double salt of the 
composition 

(NO) 3 0 3 Co'" + 3 NO • OK, 

with varying quantities of water, then separates also in the form 
of a yellow crystalline precipitate. This compound is only slightly 
soluble in water, but quite insoluble in water containing a small 
quantity of potassium nitrite or some other potassium salt. The 
corresponding nickel compound is easily soluble in water, and 
potassium nitrite is therefore used to separate cobalt and nickel. 

Cobaltous Phosphate is thrown down as a bright-red precipitate 
on adding a solution of sodium phosphate to a cobaltous salt. 

Cobaltous Carbonate. —When sodium carbonate is mixed with 


487 


Compounds of Cobalt . 

a cobaltous salt a pink precipitate is produced. This is not 
the normal carbonate, but a basic salt of the composition : 
2 CO • 0,Co + 3Co(OH) 2 + H 2 0. 

(CO 

Cobaltous Oxalate : -j ^q0 2 Co, is slightly soluble in water, and 

separates when oxalic acid is added to a cobaltous salt as a pink 
crystalline powder. 

Cobaltous Chloride : CoCl 2 + 6 H..O, crystallizes on evaporating 
its aqueous solution in fine red prisms. When heated the salt 
loses water, and becomes of a blue colour. The same change is 
produced by concentrated hydrochloric acid. It forms a crystalline 
double salt with ammonium chloride of the composition : CoCl 2 + 
NH 4 C1 + 6H 2 0. 

Cobaltic chloride, which in the free state is so unstable, unites 
with ammonia and forms a number of stable and most remarkable 
compounds, whose constitution still remains an unsolved problem. 
If an aqueous solution of cobaltous chloride is mixed with an 
excess of strong ammonia, the clear liquid exposed to the air until 
its brown colour has changed to a red, and then mixed with con¬ 
centrated cold hydrochloric acid, a brick red crystalline powder of 
the composition : 

CoC1 3 ,5NH 3 + H 2 0, 

and called Roseocobaltic chloride , separates out. The cobaltous 
chloride by standing in the air has become oxidized to a compound 
of triad cobalt. If the solution is precipitated with sulphuric acid, 
instead of with hydrochloric acid, a bright red crystalline pre¬ 
cipitate of roseocobaltic sulphate is produced. 

The dark red solution of roseocobaltic chloride becomes of a 
purple colour when boiled, , and then deposits small, lustrous, purple 
crystals of Purpureocobaltic chloride : CoC1 3 ,5NH 3 , on cooling. 
This compound differs only from roseocobaltic chloride by contain¬ 
ing one molecule of water less. 

With the formation of the purpureocobaltic chloride, a second 
compound called Luteocobaliic chloride : CoC1 3 ,6NH 3 , is also pro¬ 
duced, which is difficultly soluble in cold water, and is deposited 
when the hot solution cools as dark yellow crystals. 

If the aqueous solutions of these chlorides are digested with 
silver oxide, silver chloride is produced, together with a strongly 
alkaline liquid, from which, however, the bases cannot be separated 


488 Text-Book of Inorganic Chemistry . 

as they decompose, with evolution of ammonia, when the solutions 
are concentrated. In what form the five or six molecules of ammonia 
exist in these compounds it is still uncertain. But the constitution 
of these compounds is certainly not that expressed by the above 
empirical formulas, which represent them as double compounds of 
cobaltic chloride and ammonia. 

Cobaltous Cya7iide\ CoCy 2 .—If potassium cyanide is added to 
a cobaltous salt, a dirty-red precipitate of cobaltous cyanide is 
thrown down. This salt is insoluble in water, but readily dissolves 
in an excess of potassium cyanide, forming a double salt : 2 KCy, 
CoCy 2 , which is decomposed by hydrochloric acid. If, however, 
the clear solution of this double salt is boiled with an excess of 
potassium cyanide, hydrogen is evolved, the cobalt oxidized, and 

Potassium Cobalticyanide : KgCo^'Cy,. = 3KCy,CoCy 3 , pro¬ 
duced according to the equation : — 

2 KCy,Co 'Cy 2 + 2 KCy + H 2 0 = 3 KCy,CoCy 3 + KOH + H. 

The salt is deposited on afterwards evaporating the solution in 
yellow rhombic crystals of the above composition and without 
water. It is easily soluble in water, acids do not precipitate 
cobaltic cyanide, and it is generally similar to potassium ferri- 
cyanide, with which it is isomorphous. In the same way, the 
cobalt which it contains is neither precipitated with caustic soda 
nor with ammonium sulphide, but hydrochloric or sulphuric acid 
when added to its concentrated aqueous solution separates hydro- 
cobalticyanic acid : H 3 CoCy 6 , which separates out as white needles 
on evaporation and addition of alcohol. In solutions of the salts 
of the heavy metals, potassium cobalticyanide produces insoluble 
precipitates of its various salts. The cupric compound, for ex¬ 
ample, is a pale-blue precipitate. 

Cobalt unites with sulphur in several proportions, and produces 
the following compounds. 

Cobaltous Sulphide : CoS, is obtained as a black amorphous 
precipitate when ammonium sulphide is added to a cobaltous salt, 
and is insoluble both in water and in dilute acids. The com¬ 
pound may also be obtained as a grey compact mass by heating 
a mixture of cobalt and sulphur. 

Tricobaltic Teirasulphide : Co 3 S 4 = CoS,Co 2 S 3 , occurs in nature 


Compounds of Cobalt. 489 

in steel-grey octahedra of the regular system as the mineral 
linnceite. 

Cobalt Disulphide : CoS 2 , is not known in the free state, but 
only in combination with cobalt diarsenide as the mineral glance 
cobalt. 

Cobalt Diarsenide : Co As.,, and Cobalt Triarsenide : CoAs 3 , both 
occur in the mineral kingdom, the former as speisscobalt and the 
latter as skutterudite. 

The chief application of cobalt in the arts is for the preparation 
of the stable blue pigment known as smalt. 

Smalt is a potash glass which contains cobalt oxide instead of 
lime. It is prepared by fusing the partially roasted cobalt ores 
(zaffre ) with quartz and potassium carbonate. The potassium 
silicate then unites with the cobalt silicate, producing a dark-blue 
glass. At the same time the compounds of sulphur and arsenic 
with the other metals (especially nickel) accompanying the cobalt 
collect at the base of the crucible, producing a speiss which is 
worked for nickel. 

The blue cobalt glass is ground to an impalpable powder, and 
used in this form as a pigment. Various shades of blue are pro¬ 
duced, depending chiefly on the amount of cobalt introduced into 
the glass, which seldom exceeds 7 or 8 per cent. 

Smalt is one of the most beautiful of blue pigments, and is, at 
the same time, the most stable. It is not bleached in light, nor is 
it attacked by acids or alkalies. It is not, however, available for 
all purposes where a blue pigment is required, and since the intro¬ 
duction of so much cheap ultramarine the demand for the more 
expensive but more stable smalt has considerably diminished. 

Detection of Cobalt Compounds. 

The intense blue colour which traces of any cobalt compound 
impart to a bead of borax, in common with other kinds of glass, 
and the fact that the blue colour is permanent both in the reduc¬ 
ing and in the oxidizing flame, makes it easy to detect these com¬ 
pounds. The salts may also be recognized by their red colour 
when hydrated and blue colour when anhydrous. Sulphuretted 
hydrogen produces no precipitate in acid solutions of cobalt salts, 
and in this way cobalt may be easily separated from copper, lead, 
and all those metals whose sulphides are precipitated from acid 
liquids by sulphuretted hydrogen. Cobaltous sulphide, like the 


490 Text-Book of Inorganic Chemistry. 

corresponding sulphide of nickel, after precipitation with am¬ 
monium sulphide, does not redissolve in dilute hydrochloric acid, 
and may thus be separated from the sulphides of iron, manganese 
and zinc, which are easily soluble in dilute acids. 


NICKEL. 

Chemical Symbol'. Ni .—Atomic Weight'. 58 ’ 5 . 

This metal, like cobalt, is only found native in metallic meteor¬ 
ites, but always in larger quantities than cobalt. The chief com¬ 
pound in which nickel occurs in nature is Kupfernickel or niccolite : 
NiAs, so called on account of its reddish colour. The metal is 
prepared from this and. from cobalt speiss by removing the cobalt 
by a series of processes and then reducing with charcoal. 

Metallic nickel was formerly almost exclusively brought into 
trade as small cubes, obtained by heating the compressed oxide in 
charcoal ; it is now largely used in the form of rolled plates for 
the processes of nickel-plating. This commercial metal is far from 
pure, and usually contains copper and iron as well as arsenic and 
-cobalt. A purer metal may be obtained by heating the carbonate 
or oxide in a stream of hydrogen, and then fusing the spongy mass 
so produced under borax at a high temperature. 

Large deposits of nickel, in the form of a hydrated double sili¬ 
cate of nickel and magnesium, have recently been discovered in 
New Caledonia, and considerable quantities of the metal are now 
obtained from this source. 

Nickel is a white metal with a yellowish tinge and nearly 9*0 
specific gravity. It is malleable and ductile and, like cobalt, faintly 
magnetic. Nickel is very difficultly fusible, and only changes 
slightly in moist air; when heated in the air it becomes covered 
with a thin layer of oxide. Hydrochloric and dilute sulphuric acid 
dissolve the metal slowly, with evolution of hydrogen ; nitric acid, 
when not too concentrated, converts it into the nitrate. 

Just as the hydrated cobalt salts nearly all possess a red colour, 
the nickel salts are distinguished by their green colour. Nickel 
forms two oxides : nickelous and nickelic oxide, corresponding to 
the two oxides of cobalt. 

Nickelous Oxide : NiO, is a dirty-green powder obtained by 



Nickel. 


491 

heating its hydrate or carbonate. It dissolves in acids and forms 
green solutions. 

Nickelous Hydrate : Ni(OH) 2 , is thrown down as an apple-green 
precipitate when caustic soda is added to a nickel salt. It dries to 
a dark-green powder. 

The same precipitate is produced by ammonia, but dissolves in 
excess of the reagent to a blue liquid. 

Nickelic Oxide : Ni 2 0 3 , is a black powder produced on care¬ 
fully heating nickel nitrate. The corresponding hydrate Ni(OH 3 ) 
is obtained as a black amorphous powder when a salt of nickel is 
precipitated with sodium hypochlorite. Both nickelic oxide and 
hydrate dissolve in hydrochloric acid with evolution of chlorine and 
formation of nickel chloride ; and both break up when heated into 
nickelous oxide and oxygen. 

Nickel Sulphate: S0 2 • 0 2 Ni + 7H 2 0, crystallizes from its 
solutions in emerald-green prisms. Like other sulphates of the 
same composition, it forms double salts with the alkaline sulphates, 
of which the ammonium compound: S0 2 *0 2 Ni + S0 2 (0NH 4 ) 2 + 
6H.,0, is largely used in the arts for nickel-plating. 

Nickel Nitrate : (N0 2 ) 2 0 2 Ni + 6H 2 0, forms long green prisms, 
easily soluble in water, and deliquescing in the air. 

Nickel Carbonate. —A basic carbonate is thrown down as a green 
precipitate when sodium carbonate is added to a nickel salt. The 
precipitate, like other insoluble nickel compounds, dissolves in 
ammonia or ammonium carbonate to a blue liquid. 

Nickel Chloride : NiCl 2 + 6H 2 0, is deposited from its aqueous 
solution when strongly concentrated as small green crystals, which, 
when heated, lose water and become yellow. 

Nickel Cyanide : NiCy 2 , is thrown down as a dirty green pre¬ 
cipitate when a solution of potassium cyanide is carefully added to 
a salt of nickel. The salt is insoluble in water, but dissolves easily 
in an excess of potassium cyanide to form a double compound of 
the composition : 2 KCy,NiCy 2 + H 2 0, which crystallizes in yellow 
prisms when the solution is evaporated down. Nickel cyanide 
produces similar double salts with the cyanides of sodium, barium, 
calcium, &c., from all of which it is again precipitated by hydro¬ 
chloric acid. Potassium-nickel cyanide is not changed when boiled 


492 Text-Book of Inorganic Chemistry. 

with an excess of potassium cyanide ; no compound corresponding 
to potassium cobalticyanide is produced, nor does such a compound 
appear to exist. 

Nickel Sulphide : NiS, is found in the mineral kingdom in fine 
hair-like crystals as the mineral millerite. The same compound 
is produced as a black precipitate by adding sodium sulphide to a 
solution of a nickel salt. If precipitated with ammonium sulphide 
a portion of the precipitate dissolves and forms a dark brown 
solution. 

Nickel Alloys. 

The chief use of nickel in the arts is for the manufacture of 
various alloys, the most important of which is that known as 
German-silver, argentcm, or nickel-silver, and which contains copper, 
zinc, and nickel in varying proportions. If the quantity of nickel 
is considerable the alloy has a white colour, with a faint yellowish 
tinge. German-silver is used for a variety of purposes, and forms 
the basis of the better kind of silver-plated ware. The five- and 
ten-pfennig pieces of the Germans, corresponding to our half¬ 
pence and pence, consist of an alloy of three parts of copper and 
one part of nickel. Alloys of nickel are also used for coinage pur¬ 
poses in Switzerland, Belgium, America, and Jamaica. 

In recent years considerable quantities of nickel have been 
used for the process of nickel-plating, which consists in coating 
metallic articles with a thin layer of nickel by an electric current. 
The coating of nickel much improves the appearance, and at the 
same time only rusts slowly in the air. Nickel-plating is especially 
used for articles made of iron, which, when well deposited, it effec¬ 
tually protects from rust. 

Detection of Nickel Compounds. 

Nickel oxide dissolves in molten glass and borax, like cobalt 
oxide, but the colour is far less characteristic. In the oxidizing 
flame the borax bead is of a yellowish-brown colour, which becomes 
grey and opaque in the reducing flame from finely-divided metallic 
nickel. 

A solution of a nickel compound is characterized by the black 
precipitate of nickel sulphide produced on the addition of ammo¬ 
nium sulphide, and especially that the precipitate is slightly 
soluble, forming a dirty-brown solution. In other respects nickel 
compounds closely resemble those of cobalt, except in their colour. 


Zinc . 


493 


The salts of the two metals may be separated by the addition of 
acetic acid and potassium nitrite, when the cobalt is completely 
precipitated but the nickel remains in solution (p. 486 ). Salts of 
calcium, barium, and strontium must not be present, otherwise 
some of the nickel is precipitated with the cobalt. 


ZINC. 

Chemical Symbol : Zn.— Atomic Weight : 65 . 

Zinc is never found free in nature, but chiefly occurs in com¬ 
bination with carbonic acid as smithsonite or calamine (zinc 
carbonate: C0*0. 2 Zn), and united with sulphur as zinc blende 
(zinc sulphide : ZnS). A further important ore of zinc is a silicate 
called siliceous calamine or willemite. Zinc is volatile at a bright 
red-heat, and its oxide cannot therefore be reduced, like iron and 
other metals, with charcoal in an ordinary furnace. All the zinc 
would then be volatilized and burnt. The zinc oxide, obtained by 
roasting any of the above ores in the air, is therefore heated with 
charcoal in specially constructed vessels of fireclay, resembling 
retorts, and in the cooler parts of which the volatilized zinc 
condenses. 

Commercial zinc is never pure ; even when freed from other 
less volatile metals by repeated distillation, it always contains 
more or less arsenic. This impurity may, however, be removed 
by stirring small quantities of nitre into the fused metal; a con¬ 
siderable quantity of the zinc is then oxidized together with the 
arsenic, which is converted into potassium arsenate. If all the 
arsenic has not been removed during the first operation, the pro¬ 
cess must be repeated. The presence of arsenic in zinc is best 
detected by the aid of Marsh’s test (p. 239 ). 

Zinc possesses very remarkable physical properties, which are 
not found in the same degree in any other metal. Under some 
circumstances it is tough and hard, under others soft and malleable, 
or it may be so brittle as to be easily powdered, and finally it is 
known as a gas. Zinc acquires these very various properties at 
different temperatures. The bluish-white metal, with a lamellar, 
crystalline fracture and a specific gravity of 7 * 15 , is hard and 
tough at the ordinary temperature, and can neither be hammered 
nor rolled into plates ; for this reason zinc was for a long time only 




494 Text-Book of Inorganic Chemistry . 

used for the preparation of brass and other alloys. The metal 
became, however, much more valuable when it was discovered 
that at the boiling-point of water— i.e. between ioo° and 150 °—it 
becomes soft and malleable, and can be hammered or rolled to 
thin plates. At a higher temperature, 200 ° and above, it is as 
brittle as antimony, and can be rubbed to powder in a hot mortar. 
Zinc melts at about 412 0 , and boils at a bright-red heat— i.e. about 
1 ,ooo°. 

Zinc, when heated in a crucible to above its melting-point and 
then exposed to the air by opening the crucible, burns with a 
bright bluish flame, forming zinc oxide, which becomes distributed 
through the air of the room in light white flocks —tanaphilosophica 
of the alchemists. The non-volatile zinc oxide which remains 
behind in the crucible has a fine yellow colour, but becomes white 
when cold. 

Zinc oxidizes in moist air even more readily than iron, and 
becomes covered with a thin film of zinc oxide or basic zinc car¬ 
bonate. But this film forms a compact layer over the metal, and 
so protects it from further oxidation that it scarcely increases in 
weight when exposed for a long time to the air. For this reason 
zinc is used for a variety of purposes for which iron, that rusts so 
- easily and continuously, could not be employed. Large quantities 
of zinc are consumed in the process of galvanizing iron, which 
consists in coating iron plates or vessels with a layer of zinc, by 
immersing them in the molten metal. And one of the most 
important and useful alloys—brass—contains a large proportion 
of zinc. 

Zinc is easily dissolved by acids : hydrochloric and dilute 
sulphuric acid evolve hydrogen when acting on the metal; nitric 
acid produces nitric oxide, or if dilute, ammonia, and oxyammonia 
when sulphuric acid is also present (p. 188 ). Zinc, like iron, easily 
decomposes water at a low red heat. And, since zinc oxide forms 
compounds with the alkalies which are soluble in water, the metal 
also dissolves in caustic alkalies, and in ammonia when warmed, 
and then liberates hydrogen. 

Zinc only forms a single compound with oxygen, sulphur, or the 
halogens, respectively ; it is a dyad in all its compounds. 


Zinc Oxide : ZnO.—This compound, the only oxide of zinc, is a 
soft white powder,infusible at a red heat. It is obtained by burning 
zinc in the air or by heating its hydrate, carbonate, or nitrate 


Compounds of Zinc. 495 

When heated it becomes of a fine yellow colour, but is again white 
when cold. 

Zinc oxide is used in medicine for some purposes, but is chiefly 
employed as a white pigment in the place of white-lead. This zinc- 
white does not cover so well as white-lead, but possesses the 
advantage of not blackening in the air. Sulphuretted hydrogen, 
which is contained in traces in the air of inhabited places, gradually 
converts both the zinc and lead compounds into sulphides, of which 
lead sulphide is black, but zinc sulphide is white, and is therefore 
not noticed. 

Zinc Hydrate : Zn(OH) 2 , is obtained as a gelatinous, white 
precipitate when a small quantity of caustic soda is added to a 
solution of a zinc compound. When washed and dried it remains 
as a white powder. The hydrate dissolves in excess of caustic 
soda and produces a soluble salt, sodium zincate , in which zinc 
oxide plays the part of an acid. It also dissolves in ammonia. 

Zinc Sulphate, white vitriol'. S0 2 • 0 2 Zn + 7H 2 0, is prepared 
by dissolving zinc or zinc oxide in dilute sulphuric acid, or on the 
large scale by gently roasting blende (zinc sulphide) and extracting 
with water. It crystallizes in large rhombic prisms, and when 
heated melts in its water of crystallization. This fused mass is. 

' sometimes cast into moulds, broken into pieces when cold, and so 
brought into the market for dyeing purposes. Zinc sulphate forms 
a basic sulphate when its solution is heated with zinc hydrate ; it 
unites with alkaline sulphates to form crystalline double salts— e.g. 
SO a (OK) 2 + SO, • 0 2 Zn + 6H 2 0. 

Zinc Nitrate : (N0 2 ) 2 0 2 Zn + 6H 2 0, crystallizes from its concen¬ 
trated aqueous solution in four-sided prisms. It is very easily 
soluble in water and deliquesces in the air; when heated it melts 
in its water of crystallization. 

Zinc Carbonate : C0*0 2 Zn, occurs in nature as the mineral 
calamine and occasionally in rhombohedra isomorphous with 
calcite, as zinc-spar. The precipitate obtained when sodium 
carbonate is added to a solution of a zinc salt is always a basic 
salt, or, in other words, a compound of zinc carbonate and hydrate. 

Zinc Chloride : ZnCl 2 .—This salt may be obtained either by 
heating zinc in chlorine, or by evaporating an aqueous solution and 



496 Text-Book of Inorganic Chemistry. 

distilling the residue. The distillate obtained by either of these 
processes is a white translucent mass. It eagerly attracts water 
from the air and deliquesces and dissolves readily both in water and 
alcohol. Zinc chloride melts at 260 ° and does not appreciably 
volatilize at 400 ° ; it is therefore often used instead of an oil-bath 
to heat substances to a high and constant temperature. Its boil¬ 
ing-point is about 700 °. 

An aqueous solution of zinc chloride is obtained by dissolving 
zinc in hydrochloric acid. If this solution is evaporated, a part of 
the salt is decomposed by the water into hydrochloric acid, which 
passes away, and zinc oxide, which unites with the rest of the zinc 
chloride to form a basic salt. The concentrated solution of zinc 
chloride is powerfully caustic and cannot be filtered through paper. 
It also possesses strong antiseptic properties, and is sometimes used 
to preserve wood from decay. If a concentrated solution of zinc 
chloride is made into a paste with zinc oxide, the mixture becomes 
warm and soon sets to a snow-white hard mass of an oxy-chloride. 
If finely powdered glass is added to the mixture, a very hard 
cement is obtained ; such mixtures are employed as a cheap stopping 
for decayed teeth. 

Zinc Sulphide : ZnS.—The zinc sulphide occurring in nature as 
blende is of a yellow or brown colour, but that obtained artificially 
by precipitating a solution of a zinc salt with ammonium sulphide 
is a pure white powder. Hydrochloric acid readily dissolves it, 
with evolution of sulphuretted hydrogen and formation of zinc 
chloride. Zinc sulphide is not precipitated when sulphuretted 
hydrogen is led into a solution of a zinc salt acidulated with 
hydrochloric or sulphuric acid ; but the gas throws down all the 
zinc from a solution of its acetate. 


Detection of Zinc Compounds. 

If any zinc compound is mixed with sodium carbonate and 
heated on a piece of charcoal in the inner blowpipe flame, the zinc 
is reduced, volatilized and burnt, and produces an incrustation on 
the charcoal which is yellow when hot, and white when cold. 
Zinc compounds are further characterized by the white precipi¬ 
tate which they give with ammonium sulphide. Zinc may be 
easily separated from iron, nickel, and cobalt by the addition of 
caustic soda in excess. Zinc oxide then alone dissolves, and can 


Cadmium. 


49 7 


afterwards be precipitated as white zinc sulphide by passing sulphu¬ 
retted hydrogen through the alkaline solution. Zinc hydrate and 
carbonate dissolve also in caustic potash, in ammonia, or in am¬ 
monium carbonate, and produce soluble double compounds. 


CADMIUM. 

Chemical Symbol : Cd.— A tomic Weight'. 112 . 

Like indium, cadmium is nearly always found associated with 
zinc in nature, but zinc ores usually contain far more cadmium 
than indium—often as much as five per cent. Cadmium is ob¬ 
tained as a bye-product in the extraction of zinc, and as it is more 
volatile than this metal, is found chiefly in the first portions which 
come over when the zinc is distilled. Cadmium sulphide is occa¬ 
sionally found as the rare mineral, greenockile. 

Cadmium is a tin-white metal, of specific gravity 8 - 6 , melting 
at 315 0 and boiling at about 770 °. Unlike zinc, it is soft at the 
ordinary temperature, and can be rolled into sheets or drawn into 
wire. The metal remains unaltered in the air, but burns brilliantly 
at a red-heat, producing brown vapours of cadmium oxide. It 
dissolves slowly in hydrochloric or dilute sulphuric acid, with evo¬ 
lution of hydrogen ; nitric acid dissolves it more readily. 

All compounds of cadmium contain the metal in the form of a 
dyad, in which respect it resembles zinc. 

Cadmium Oxide : CdO, is obtained as a brown amorphous 
powder, soluble in acids, when the metal is heated in the air or 
when the nitrate is glowed. Cadmium Hydrate'. Cd(OH) 2 , is 
thrown down as a white precipitate on adding caustic soda to a 
solution of a cadmium salt. The precipitate is insoluble in excess 
of caustic soda, and so differs from zinc hydrate; it is, however, 
like the zinc compound, easily soluble in ammonia. 

The cadmium salts, with the exception of the sulphide, are all 
colourless. 

Cadmium Sulphate : 3S0 2 0 2 Cd 4 8H 2 0, forms large, colourless, 
easily soluble crystals. Cadmium Nitrate'. (N0 2 ) 2 0 2 Cd 4H 2 0, crys¬ 
tallizes in deliquescent prisms. Cadmium Carbonate : CO • 0 2 Cd, is 
precipitated as a white powder by adding sodium carbonate to a 
solution of a cadmium salt. Cadmium Chloride : CdCl 2 + 2H 2 0, 

K K 



498 Text-Book of Inorganic Chemistry. 

crystallizes in colourless prisms, which effloresce when exposed to 
the air and are easily soluble in water. Cadmium Sulphide : CdS, 
is thrown down when sulphuretted hydrogen is led into a solution 
of a cadmium salt as a bright yellow precipitate, which is insoluble 
in dilute, but dissolves in concentrated hydrochloric acid. It is 
distinguished from arsenious sulphide, which has a similar colour, 
by its insolubility in alkaline sulphides, including ammonium sul¬ 
phide. Cadmium sulphide is used as a stable yellow pigment. 


Detection of Cadmium Compounds. 

Compounds of cadmium when heated on charcoal in the 
inner flame of the blowpipe are reduced to the metallic state, 
but the metal at once volatilizes and burns in the outer flame, pro¬ 
ducing a brown incrustation of cadmium oxide on the charcoal. 

When a current of sulphuretted hydrogen is let through an 
acid solution of a cadmium salt, a yellow precipitate of cadmium 
sulphide is produced, insoluble in alkaline sulphides (distinction 
from arsenic). This reaction serves to separate cadmium from 
zinc, which is not precipitated from its acid solutions by sulphu¬ 
retted hydrogen. If ammonia is added to a solution of a cadmium 
salt, a white precipitate of cadmium hydrate is formed, which dis¬ 
solves in an excess of the reagent. Caustic soda produces the 
same precipitate, but does not redissolve it when an excess is 
added. Cadmium may be separated from copper by boiling the 
mixed sulphides with dilute sulphuric acid ; the cadmium sulphide 
then dissolves and the cupric sulphide remains behind. 


LEAD. 

Chemical Symbol : Pb.— Atomic Weight'. 207 . 

The most important form in which lead occurs in nature is as 
the sulphide : PbS—the mineral galena. Lead is also found in 
smaller quantities as the sulphate— anglesite , as the carbonate 
cerussite , and less commonly as the chromate, phosphate, and 
molybdate. 

Two processes are chiefly used for the extraction of lead from 
galena accordingly as the ore is more or less pure. The impure 



Lead, 


499 


ores are either roasted and reduced with carbon or else simply 
smelted with iron. In the latter case the iron unites with the 
sulphur for which it has a stronger attraction than the lead, and 
forms a fusible slag of ferrous sulphide with the impurities, which 
covers the heavier reduced lead. In the second method-the so- 
called air reduction process —the sulphur is indirectly oxidized by 
the oxygen of the air. The ore is first carefully roasted in a 
specially constructed furnace, so that one portion is converted into 
lead sulphate, another portion into lead oxide, and the rest remains 
as unchanged lead sulphide. The whole is afterwards well mixed 
and heated more strongly with exclusion of air, when the sulphur in 
t e unchanged sulphide is burnt by the oxygen contained in the 
lead oxide and sulphate, and the reduced lead runs from the hearth 
of the furnace into moulds 

PbS + 2 PbO = 3 Pb + S0 0 . 

PbS + S0 2 -0,Pb = 2 Pb + 2 s6 3 . 

Lead when thus smelted often contains silver. If the quantity 
is sufficient to pay for its extraction, 1 the lead is converted by 
oxidation into lead oxide, the silver remaining unchanged. This 
process, called cupellation , is performed in an oxidizing reverbera¬ 
tory furnace, of which the German form is shown in fig. 66. This 



Fig. 66. 


furnace consists of a porous clay hearth, on which the argentiferous 
lead, a, is placed, and which is kept fused by the heated gases 

1 I.e. about one-tenth per cent. 


K K 2 





500 Text-Book of Inorganic Chemistry. 

from the fire, F, while a continuous blast of air is blown through the 
openings a , a , over the surface of the metal in order to oxidize the 
lead to lead oxide. The lead oxide melts as it is produced, and is 
partly absorbed by the porous clay, but mostly flows away by 
side openings into vessels provided to receive it. The silver, 
however, remains unoxidized, and the lead oxide is therefore free 
from this metal, except towards the end of the operation. Finally, 
when all the lead has been converted into lead oxide, and the 
silver is coated with a thin film of the oxide, beautiful iridescent 
tints are observed on its surface, due to the colours of thin plates ; 
immediately afterwards the bare surface of the molten silver flashes 
out, and the operation is over. In the English form of cupel 
furnace, the hearth is made of bone ash, and the cover is not 
moveable. The lead oxide (litharge), of which large quantities 
are produced by cupellation, is afterwards again converted into 
metallic lead by reduction with charcoal or coal. 

Two other processes are now largely used for extracting silver 
from lead, when the quantity is insufficient to pay the expense of 
cupelling. The first of these, known as Pattinson's pf'ocess , depends 
upon the fact that when fused argentiferous lead is allowed to cool 
slowly, the crystals of lead which first separate out are almost 
entirely free from silver ; and if these crystals are removed and 
the operation repeated several times, a sample of lead is at last 
obtained which is rich enough in silver to pay for cupelling. A 
second method, called Parkes ’ process , is based upon the property 
of zinc to form an alloy with silver, but not with lead. If a small 
quantity of zinc is added to the molten argentiferous lead, and the 
mass allowed to cool, the alloy of zinc and silver rises to the sur¬ 
face and can be removed during solidification. The zinc can then 
be abstracted from this alloy by dissolving it in dilute sulphuric 
acid, or by distilling it off. 

Lead is a lustrous metal of bluish-grey colour, and soft enough 
to be cut with a knife, and to mark paper. It is malleable and 
ductile, but possesses little tenacity. The metal has a specific 
gravity of 11 * 4 , melts at 32 5 0 , and volatilizes at a white heat, though 
it cannot be distilled. When exposed to moist air, it becomes 
dull and coated with a thin grey layer of lead carbonate. Pure 
water containing air also attacks it and dissolves a small quantity 
of the lead. But if the water contains traces of certain salts dis¬ 
solved in it— e.g. calcium sulphate—the leaden pipes through 
which the water passes are not attacked. All lead compounds are 


Oxides of Lead. 501 

poisonous, and especially when dissolved in water, and hence 
drinking water containing only traces of dissolved lead ^probably 
as its acid carbonate) is injurious to health. The presence of lead 
in a sample of drinking water may be readily proved by the addition 
of sulphuretted hydrogen water, or by leading a stream of the gas 
through it, when if lead is present the water becomes of a brown 
colour from the lead sulphide formed. 

Hydrochloric acid and dilute sulphuric acid scarcely act upon 
lead, either in the cold or when warmed. Hot concentrated sul¬ 
phuric acid dissolves lead slightly, but only so slightly that sul¬ 
phuric acid may be evaporated to a certain strength in leaden 
dishes (see p. 159). The best solvent for lead is nitric acid, which 
dissolves it readily, producing lead nitrate and nitric peroxide. 
From this salt, or from a solution of lead acetate, zinc precipitates 
metallic lead. The lead is then deposited in lustrous crystalline 
plates, and, as the metal in this form possesses some similarity 
with leaves and boughs, it is sometimes called a ‘ lead-tree.’ 

Lead unites with oxygen (not with sulphur or the halogens) in 
three proportions, and produces the compounds— lead suboxide : 
Pb 0 0 , lead oxide : PbO, and lead peroxide : PbO, 2 . A fourth oxide 
is red-lead : Pb 3 0 4 or Pb 4 0 5 , which, however, must be considered 
as a compound of lead oxide and lead peroxide. 

lead Suboxide : Pb o 0 .—This compound, which is only of 
scientific interest, is a dark-grey indifferent powder, and is pro¬ 
duced when lead oxalate is heated to 300° in an oil-bath as long as 
gas is given off. The decomposition which the oxalate then under¬ 
goes is represented in the equation :— 

2 {co°* Pb = pb *° + 3C ° 2 + ca 

Lead suboxide, when heated in the air, is converted into lead 
oxide ; it does not unite with acids, but is decomposed by them 
into lead oxide and lead. 

lead Oxide, Litharge, Massicot : PbO, occurs as a yellow or 
reddish powder, or as crystalline scales, according to the method 
used for its preparation. It easily melts, and solidifies on cooling 
to a crystalline mass. Lead oxide may be obtained by heating 
either lead hydrate, carbonate, or nitrate, or by the oxidation of 
molten lead in the air. Most of the litharge brought into trade is 
derived from the cupellation of argentiferous lead. 


502 Text-Book of Inorganic Chemistry. 

Lead oxide is insoluble in water, but dissolves easily in nitric 
or acetic acid. It is a powerful base, and as most of its salts are 
insoluble in water, a solution of the soluble acetate is used in the 
laboratory for the detection of acids. 

Litharge is largely used in the arts : considerable quantities are 
consumed in the manufacture of flint-glass, and as a cheap glaze 
for stone and earthenware. Nearly all other lead compounds are 
prepared from it— e.g. lead acetate and nitrate, red-lead, lead- 
plaster, &c. 

lead Hydrate : Pb(OH) 2 , is produced as a white, voluminous 
precipitate, when a solution of lead acetate is mixed with a large 
excess of caustic soda, until the solution is strongly alkaline, and 
until a portion of the hydrate has been redissolved by the soda. If 
a smaller quantity of caustic soda is used than is required to com¬ 
pletely precipitate the lead, the lead hydrate is mixed with basic 
lead acetate ; and as sodium acetate itself has an alkaline reaction, 
■the liquid must be strongly alkaline before the soda is in excess. 
Lead hydrate preserves its white colour when washed and dried. 
It is quite insoluble in water, but unites with acids more readily 
than lead oxide, and even attracts carbonic acid from the air. 
When heated it loses water and yields lead oxide. 

Red-lead : Pb 3 0 4 = 2PbO,PbO a .—If litharge (lead oxide) is 
heated in the air to about 400°, but not much above, it absorbs 
oxygen, and is converted into a bright brick-red crystalline powder. 
This substance is called red-lead or minium , and is used both as a 
pigment and for a number of other technical purposes. This 
oxide cannot unite with acids and form salts like lead oxide ; and 
although dilute nitric acid acts upon it even in the cold, it only 
abstracts lead oxide from it, and leaves brown lead peroxide be¬ 
hind. Red-lead evolves chlorine when warmed with hydrochloric 
acid, owing to the lead peroxide which it contains. When strongly 
heated it loses oxygen and is converted into yellow lead oxide. 

X*ead Peroxide : Pb 0 2 .—The preparation of this substance 
from red-lead by the action of dilute nitric acid has just been de¬ 
scribed. When repeatedly boiled and washed with water and dried, 
it remains as a brown amorphous powder, with a purple tinge— 
whence it is sometimes called the puce-coloured oxide of lead. 
Lead peroxide may also be obtained by precipitating a solution of 


503 


Salts of Lead. 

lead nitrate with sodium hypochlorite. No solvent is known for 
lead peroxide ; like manganese peroxide, which it resembles in 
other points, it is not attacked by nitric acid. When warmed with 
hydrochloric acid it evolves chlorine ; and when heated alone oi 
with sulphuric acid oxygen is given off, and in the latter case, white 
insoluble lead sulphate remains behind. It unites with dry sul¬ 
phurous anhydride, with considerable evolution of light and heat, 
and produces the same substance (p. 152). If lead peroxide is 
warmed with nitric acid, and finely powdered sugar gradually added, 
the sugar is oxidized to carbonic acid and water by a half of the 
oxygen contained in the peroxide, and a clear solution of lead 
nitrate is soon obtained. 

Salts of Lead. —All the salts of lead contain the metal as a 
dyad, and correspond to lead oxide; none have yet been prepared 
containing lead in the tetrad form, although it is possible that a 
tetrachloride corresponding to the peroxide might be produced by 
the action of hydrochloric acid on the peroxide ; but if produced it 
is so unstable that its isolation is exceedingly difficult. 

Most of the lead salts are insoluble in water. Those which 
dissolve in water include the nitrate and chlorate, as well as 
the acetate. And although the acetate belongs strictly to organic 
chemistry, it is the most important soluble lead salt, and may 
well be included with the inorganic compounds. Many salts of 
lead are largely used both in the arts and in medicine. All lead 
compounds, and especially those which are soluble in water, are 
very poisonous. The soluble compounds are characterized by 
their sweet taste. 

lead Sulphate : S 0 2 - 0 . 2 Pb, occurs in nature as the mineral 
anglesite , and may be prepared artificially by precipitating a soluble 
lead salt with dilute sulphuric acid or with a solution of a sulphate. 
It is only very slightly soluble in water and dilute acids, but hot 
concentrated sulphuric acid dissolves not inconsiderable quantities 
of it. Strong hydrochloric acid converts it partially into lead 
chloride, with the liberation of the corresponding quantity of sul¬ 
phuric acid. Lead sulphate dissolves readily in basic ammonium 
tartrate^ for example, when digested with tartaric acid and then 
ammonia added in excess. Ammonium sulphate also forms a 
soluble crystalline double salt with it, which is, however, decom¬ 
posed by water, with separation of lead sulphate. 


504 Text-Book of Inorganic Chemistry. 

lead Nitrate : (N 0 2 ) 2 0 2 Pb, is obtained by dissolving lead or 
lead oxide in a slight excess of nitric acid, and crystallizes from 
the solution containing free acid as colourless, transparent, hard 
octahedra, without water of crystallization. It is soluble in water, 
but insoluble in strong nitric acid, by which it is precipitated from 
its concentrated aqueous solution. If the aqueous solution is 
boiled with powdered lead oxide, some of this substance dissolves 
and produces a difficultly soluble basic salt of the composition : 
N 0 2 ; 0 Pb( 0 H), which crystallizes out on cooling. Lead nitrate 
decomposes when heated into lead oxide, oxygen, and nitric oxide 

fp. 196). 

Lead Phosphate : (P 0 ) 2 0 6 Pb 3 , is thrown down as a white pre¬ 
cipitate when sodium phosphate is added to a solution of a lead 
salt. It is insoluble in water and acetic acid, and forms double 
compounds with other lead salts— e.g. the nitrate and chloride. 
The mineral pyromorphite is a double compound of lead phosphate 
and chloride: 3(PO) 2 O tJ Pb 3 + PbCl 2 , and is isomorphous with 
apatite (p. 396). 

Lead Silicates. —Silica and lead oxide may be fused together 
in almost any proportion, producing a vitreous mass of varying 
composition, which is an essential constituent of flint or lead-glass, 
and strass (p. 403). A lead silicate of definite composition has not 
yet been prepared. 

lead Carbonate : C 0 - 0 2 Pb.—The normal salt of the above 
composition is found in nature as the mineral cerussite , isomor¬ 
phous with arragonite, and may be prepared artificially as a white 
powder, insoluble in water, when a solution of lead nitrate is pre¬ 
cipitated with ammonium carbonate. 

Basic Bead Carbonate, White Bead : 2CO • 0 2 Pb + Pb(OH) 2 . 
—Far more important than the preceding compound is a basic car¬ 
bonate of lead, very largely used in the arts as a white pigment. 
W hite lead is manufactured by a variety of methods, of which the 
following two are the most important. 

The French or English process depends upon the fact that a 
stream of carbonic acid, when led into a solution of basic lead 
acetate, precipitates the whole of the lead oxide which the salt 
contains in excess of the normal compound as white lead. Finely 


White Lead. 


505 


powdered litharge (lead oxide) is boiled with dilute acetic acid 
until no more is dissolved, when the normal acetate is first obtained, 
and then the basic salt. On passing a stream of carbonic acid 
through the clear solution, a precipitate of white lead is produced 
with a solution of the normal acetate. This solution can be again 
converted into the basic acetate by boiling with litharge, and from 
this carbonic acid again throws down a further quantity of white 
lead, and so on. Thus, with a small quantity of acetic acid a large 
quantity of white lead can be obtained. 

A second and older method, called the Dutch process, because 
first carried on in Holland, consists also in decomposing a basic 
lead acetate with carbonic acid, but the arrangement by which this 
reaction is brought about is different. Plates of lead coiled into a 
spiral or cast leaden gratings are placed in earthenware pots con¬ 
taining a layer of strong vinegar at the bottom. The pots are then 
loosely closed with a leaden lid and imbedded in horse-dung or 
spent tan. After several weeks the vessels are withdrawn and 
opened, when the leaden plates are found to be strongly corroded 
and covered with a loose layer of white lead. This is removed, 
and the plates returned to the pots until they are completely 
corroded away. The process by which the white lead is produced 
is a simple one. The warmth produced by the decomposition of the 
dung or spent tan gradually evaporates the vinegar which attacks 
the lead, and. in the presence of the oxygen of the air, produces 
basic lead acetate. The acetate is then decomposed as fast as it is 
produced by the carbonic acid evolved during the putrefaction of 
the dung or spent tan, producing white lead, which remains loosely 
adhering to the leaden plates. White lead prepared in this way is 
more highly valued as a pigment than that obtained by any other 
process. The colour is purer and it covers better. 


laead Acetate, Sugar opLead : A 2 - 0 2 Pb + 3 H 2 0 = (C 2 H 3 0 ) 2 0 2 Pb 
+ 3H 2 0. a —Normal lead acetate is obtained by dissolving litharge 
in a small excess of acetic acid, and crystallizes in colourless, trans¬ 
parent, glistening prisms when the solution is evaporated down. 
Of all lead salts it is the most soluble in water; it also dissolves in 
alcohol. 

1 The abbreviated symbol A stands for the radical of acetic acid (acetyl) . 
C<jH 5 0, and has the same relation to acetic acid : A'OH, asnitryl : NOo, has to 
nitric acid : N0 2 '0H. 


506 Text-Book of Inorganic Chemistry . 

If a solution of the normal acetate is boiled with lead oxide it 
combines with this substance and produces various basic acetates, 
of which those containing the greatest quantity of the oxide are 
insoluble in water. The monobasic acetate of the composition : 
A.,0>Pb + Pb'OH), = A.OPb(OH), is easily soluble in water and 
reacts alkaline. It parts with one half ot its lead when carbonic 
acid is led through its solution. 

Iiead chromate : CrO., *0 2 Pb, is found in nature as the 
mineral crocoisite , and may be obtained as a bright yellow powder 
by precipitating a soluble lead compound with potassium chromate 
or dichromate. This precipitated lead chromate when washed 
and dried is largely used as a yellow pigment under the name of 
chrome-yellow. It melts when heated, and at higher temperatures 
evolves oxygen. If chrome-yellow is digested with caustic potash 
or soda, or if boiled with normal potassium chromate, it loses 
chromic acid and is converted into a basic lead chromate of a 
bright red colour and the composition: CrO.,-0.,Pb + PbO. This 
substance is the chrome-red of commerce, and is also largely used 
as a pigment. 

lead Chloride : PbCL>.—This salt, which is only slightly 
soluble in cold water but more in hot water, is deposited as a 
heavy white precipitate when hydrochloric acid or a soluble 
chloride is added to a solution of lead acetate. It crystallizes 
from its hot saturated solution in colourless lustrous scales or 
needles. When strongly heated, the salt melts, and solidifies on 
cooling to a horny mass which may be cut with a knife. 

Lead chloride unites with lead oxide in varying proportions 
and produces different basic chlorides or oxychlorides. One of 
these is Cassel yellow ; it is prepared by heating a mixture of 
litharge and sal-ammoniac, and is used as a figment. Another 
oxychloride, of the composition : PbCL,PbO, is the mineral mat- 
lockite. If an aqueous solution of lead chloride is boiled with lead 
hydrate, the whole of the former salt is precipitated as insoluble 
oxychloride. 

I.ead iodide : Pbl 2 , is thrown down as a pale yellow precipi¬ 
tate when potassium iodide is added to a solution of lead acetate. 
It is nearly insoluble in cold water, but boiling water dissolves it 
slightly. This hot solution is colourless, but on cooling deposits 


Alloys of Lead. 507 

the dissolved lead iodide in lustrous golden scales. Lead iodide 
forms a double compound with potassium iodide. 

Lead Cyanide : PbCy 2 , is a white flocculent precipitate obtained 
by adding potassium cyanide to lead acetate. It readily dissolves 
in an excess of potassium cyanide, and the solution then contains 
the double salt: 2KCy,PbCy 3 . 

Lead Sulphide : PbS, occurs in nature as galena, the most 
abundant ore of lead. Galena is found either in compact masses 
of a grey colour, or else crystallized as cubes, usually with sub¬ 
ordinate faces of the octahedron. 

The same compound may be obtained as a black amorphous 
precipitate by leading a stream of sulphuretted hydrogen into a 
solution of a lead salt. The precipitate retains its colour when 
dry, and if heated out of access of air, melts without decomposition; 
but if roasted in the air both sulphur and lead undergo oxidation, 
and there are produced sulphurous anhydride with lead oxide and 
sulphate. Lead sulphide is scarcely attacked by hydrochloric acid. 
Dilute nitric acid dissolves it with formation of lead nitrate and 
free sulphur. Concentrated nitric acid converts it into the white 
sulphate, as the sulphur is then oxidized, 
j Numerous double sulphides of lead and other metals, especially 
antimony, are found in nature; the minerals zinkenite, jamesonite, 
1 and boulangerite , are all double sulphides of lead and antimony in 

| varying proportions. 

Besides the salts of lead oxide in which this compound plays 
j the part of a base, others are known in which it is combined with 
strong bases as an acid. These compounds—called plmnbates — 
are produced by digesting lead hydrate with the aqueous alkalies 
or with lime- or baryta-water, in which the hydrate easily dis¬ 
solves. The salts so obtained— eg. potassium plumb ate, of the 
probable composition Pb(OK).,—have not yet been prepared pure 
enough to determine their exact composition. 

Alloys of Xiead. — Metallic leadjs not only used alone for the 
manufacture of water-pipes, for covering roofs, and for sulphuric 
acid chambers, but it is also a constituent of many important alloys. 
One of these is type-metal, an alloy of lead and antimony. German 
type-metal contains from 10 to 25 parts of antimony to 100 parts of 




508 Text-Book of Inorganic Chemistry . 

lead, while that used in England contains about 30parts of antimony 
and 30 parts of tin to 100 parts of lead, and sometimes a little 
copper as well. The larger the quantity of antimony, the harder 
and more brittle is the alloy produced. The addition of a small 
quantity of arsenic (at most o‘8 per cent.) to lead makes it suitable 
for the manufacture of shot. The alloys of lead and tin are de¬ 
scribed under the latter metal. 

Detection of Dead Compounds. 

Scarcely any other metal can be so easily detected and separated 
as lead. Any compound of lead when mixed with sodium carbonate 
and heated on charcoal in the reducing flame of the blowpipe 
yields small soft globules of metallic lead, which mark paper, 
and which are converted into yellow lead oxide in the oxidizing 
flame. In the borax bead, as in other forms of molten glass, lead 
oxide dissolves without imparting any colour. 

Sulphuretted hydrogen when passed through a neutral or acid 
solution of a lead salt precipitates black lead sulphide. The same 
solution gives a white precipitate of lead sulphate with dilute 
sulphuric acid, a white precipitate of lead chloride with hydro¬ 
chloric acid, which dissolves in hot water, and crystallizes out in 
needles on cooling, a yellow precipitate of lead iodide with potas¬ 
sium iodide, and a white precipitate of lead hydrate with caustic 
soda, soluble in excess of the reagent. Ammonia produces no pre¬ 
cipitate when added to a solution of lead acetate or only a slight 
one, after some time, but at once throws down basic lead nitrate 
when added to a solution of lead nitrate. The yellow precipitate 
which is obtained when potassium dichromate is added to a lead 
salt is also produced with a bismuth salt, but lead chromate may 
be distinguished from the bismuth salt by its solubility in caustic 
soda. 

The separation of lead from the other metals precipitated from 
their acid solutions by sulphuretted hydrogen is best effected by 
the addition of sulphuric acid and alcohol. Lead sulphate is 
slightly soluble in water, less in dilute sulphuric acid, and insoluble 
in alcohol. 



Thallium . 


509 


THALLIUM. 

Chemical Symbol : Tl .—Atomic Weight : 204. 

This remarkable element occupies an altogether anomalous 
position among other metals, for it unites in itself the chemical and 
physical properties of two groups of metals which are as different 
as possible from one another. On the one hand it has much 
similarity with lead and silver, and on the other with the alkali- 
metals, especially potassium. 

It resembles lead in possessing nearly the same atomic weight 
and specific gravity, and in its difficultly soluble chloride and 
insoluble iodide. Its sulphate is also difficultly soluble in water, 
though not to such an extent as lead sulphate ; its neutral solutions 
are precipitated by sulphuretted hydrogen, and the metal is set 
free from its soluble salts by metallic zinc. 

It differs from lead, but resembles silver and potassium, in its 
atomicity, which in the above compounds is always unity, and is 
similar to potassium in many other points. Although it does not 
decompose water like potassium, it easily dissolves in water con¬ 
taining air, and produces thallium hydrate, which is readily soluble 
in water, and of which the solution reacts strongly alkaline. Still 
more remarkable is the fact that a metal so similar to lead as 
thallium should possess a carbonate soluble in water, and even 
reacting alkaline. Twenty-five years ago, when thallium was un¬ 
known, it was universally held that only the alkali-metals yield 
soluble carbonates, and it was not thought possible that the car¬ 
bonate of a heavy metal could be soluble in water; still less was 
it imagined that the sulphate of a heavy metal could displace 
potassium sulphate in common alum without changing the physical 
properties of the salt. We now know that thallous carbonate is 
soluble in water, giving an alkaline solution, and that thallous 
sulphate can unite with aluminium sulphate and produce a true 
alum. And further, thallous chloride, just like potassium chloride, 
can unite with platinum chloride, and produce a double salt inso¬ 
luble in water. 

Thallium is a rare element, and although present in many ores 
and mineral waters, is only found in minute quantities. The 
mineral richest in thallium is crookesite (called after Crookes, the 
discoverer of the element), which contains about 18 per cent, of 


5io 


Text-Book of Inorganic Chemistry. 


the metal, but which is extremely rare. Thallium is best prepared 
from some varieties of iron or copper pyrites. When pyrites con¬ 
taining traces of the metal is used for the manufacture of sulphuric 
acid, the thallium passes over with the gases, and is then contained 
in the flue-dust or in the deposit which forms on the floor of the 
leaden chambers. This flue-dust or deposit when boiled with 
dilute sulphuric acid yields an impure solution of thallous sulphate, 
which when evaporated down and mixed with hydrochloric acid 
gives a precipitate of thallous chloride. This precipitate is filtered 
off, well washed with cold water, and then decomposed with con¬ 
centrated sulphuric acid, which converts it into thallous sul¬ 
phate. On dissolving this salt in water and dipping a plate of 
zinc (free from lead) into it, the thallium is precipitated as a 
crystalline, spongy mass, which is then pressed together and fused 
in a crucible. 

Thallium so obtained is of a tin-white colour when freshly cut; 
it is very soft, and, like lead, may be cut with a knife or even 
scratched with the nail, and easily marks paper. It has a specific 
gravity of 11*9, melts at 290°, and volatilizes at a red heat. Thal¬ 
lium readily oxidizes when exposed to the air and becomes covered 
with a thin layer of oxide, but, notwithstanding its strong attraction 
for oxygen, it does not decompose water even at the boiling tem¬ 
perature. It preserves its bright surface under water which has 
been freed from dissolved air by boiling, and is therefore best kept 
under this liquid. But if the water is exposed to the air the 
thallium gradually becomes converted into thallous hydrate and 
thallous carbonate, both of which dissolve in the water. It easily 
dissolves in dilute nitric or sulphuric acid, but less readily in 
hydrochloric acid on account of the insolubility of its chloride. 

With oxygen, sulphur, and the halogens, thallium forms two 
series of compounds—the thallous and thallic compounds. In 
the former it is a monad element, and in the latter a triad. The 
former (thallous) series of compounds are the more stable, and 
thallous oxide is a much stronger base than thallic oxide/ All 
thallium compounds are poisonous. 

If thallium is allowed to remain for a long time under water 
which has access to air or oxygen free from carbonic acid, the 
aqueous solution becomes strongly alkaline from the thallous 
hydrate : T 10 H, which has been produced. When this solution 
is evaporated down, the hydrate crystallizes out in yellow rhombic 
prisms with 1 molecule of water. The same substance may be 


Compounds of Thallium . 5 1 r 

more easily obtained by precipitating a solution of thallous sul¬ 
phate with baryta-water, filtering off the barium sulphate produced, 
and evaporating the clear solution. It is soluble both in water 
and alcohol. When heated to ioo° it decomposes into water and 
thallous oxide : Tl., 0 , which forms a black powder, melting at 
about 300° to a dark yellow liquid. The oxide attracts moisture 
from the air and is again converted into the hydrate. 


Thallous Sulphate : S 0 2 (OTl) 2 , crystallizes in colourless 

rhombic prisms isomorphous with potassium sulphate, and unites 

with aluminium sulphate to form thallium-alum : gQ 2 j- + 

( Old 

I 2 H 2 0 (p.425). An acid sulphate : S 0 2 j is also known. 


Thallous Phosphate ; PO(OTl) 3 .—If a neutral solution of a 
thallous salt is mixed with sodium phosphate and a few drops of 
ammonia added, thallous phosphate soon begins to crystallize out, 
and afterwards nearly fills the vessel. The crystals are long, 
glistening prisms, of the above composition, and almost insoluble 
in water containing a trace of ammonia. 

Thallous Carbonate : CO (0 ri) 2 , is obtained by saturating a 
solution of thallous hydrate with carbonic acid, and crystallizes in 
colourless lustrous prisms when the solution is evaporated down. 
It is tolerably easily soluble in water, and the solution reacts alka¬ 
line. It loses carbonic acid when heated, and is converted into 
thallous oxide. 

Thallous Chloride : T 1 C 1 , is thrown down as a white curdy 
precipitate when a solution of a thallous salt is mixed with hydro¬ 
chloric acid. It is scarcely soluble in cold water, but dissolves 
more readily when heated. It easily melts when heated and then 
volatilizes. Thallous chloride when mixed with platinic chloride 
forms a yellow insoluble double chloride of the composition : 
2TlCl,PtCl 4 , corresponding to the compound of potassium and 
platinic chlorides. 

Thallous iodide : Til, like lead and silver iodides, is nearly 
insoluble in water, and is produced as a yellow amorphous pre¬ 
cipitate when potassium iodide is added to a solution of a thallous 

salt. 


512 Text-Book of Inorganic Chemistry. 

Thallous Sulphide : T 1 ,S, is obtained as a black precipitate 
when excess of ammonium sulphide is added to a solution of a 
thallous salt, and is insoluble in excess of the reagent. Sulphu¬ 
retted hydrogen produces the same precipitate in a solution of 
thallous acetate. 

Thallic Compounds. —Thallic hydrate and oxide, unlike the 
corresponding thallous compounds, are insoluble in water, but are 
dissolved by acids to form soluble thallic salts. Thallic hydrate . 
TIO(OH), is produced as a brown amorphous powder when freshly 
precipitated thallous chloride is digested with an excess of sodium 
hypochlorite. At ioo° it breaks up into water and thallic oxide : 
Tl a O s , a black powder insoluble in water, but dissolving in hydro¬ 
chloric acid to form thallic chloride. When heated with concen¬ 
trated sulphuric acid, oxygen is given off and thallous sulphate 
remains behind ; when heated alone it breaks up into thallous 
oxide and oxygen. 

Thallic Sulphate : (S 0 2 ) 3 0 6 T 1 /" + 7 H. 2 0 , crystallizes in colour¬ 
less thin scales when a solution of thallic hydrate in dilute sul¬ 
phuric acid is evaporated down. It loses only 6 molecules of water 
when heated to 220°, and is decomposed by water even in the cold, 
-with separation of thallic hydrate. It forms a difficultly soluble 
double salt with potassium sulphate, which, however, is not a true 
alum. 

Thallic Chloride : T 1 C 1 3 , is produced when thallous chloride is 
gently heated in a stream of chlorine, or when thallic hydrate is 
dissolved in hydrochloric acid. On evaporating its solution it 
crystallizes in easily soluble, deliquescent prisms, which decom¬ 
pose, when heated, into thallous chloride and free chlorine. 

Thallic Sulphide : T 1 . 2 S 3 , may be obtained by fusing thallium 
with an excess of sulphur. It then forms a black soft mass, which 
easily melts, and is brittle at low temperatures. 


Detection of Thallium Compounds. 

The compounds of thallium possess such characteristic proper¬ 
ties, that it is an easy matter to detect them, and to separate them 
from the compounds of other metals. The non-precipitation of 
thallium compounds by sulphuretted hydrogen from their acid 


Bismuth. 


513 


solutions, the solubility of thallous sulphate in hot water, the in¬ 
solubility of the chloride and iodide, and the solubility of thallous 
hydrate and carbonate in water, offer the best means of detection 
and separation. 

Minute traces of thallium can be readily recognized with the 
spectroscope ; its spectrum consists of a single intense green line 
(see table). It was this property of thallium which led to its dis¬ 
covery, and from which it derived its name ( 6 d\Xos = a green 
twig). 


BISMUTH. 

Chemical Symbol : Bi .—Atomic Weight : 210. 

Bismuth is only sparsely distributed in nature, and nearly 
always occurs in the free state. From the earthy impurities-which 
accompany it, the metal is separated either by simply melting it in 
inclined iron tubes, or else by roasting in the air, and then heating 
with carbon and a slag in suitable crucibles. Commercial bismuth 
always contains traces of other metals, including arsenic. It may 
be purified by fusion with nitre, which converts the arsenic into 
potassium arsenate and oxidizes other impurities. To prepare pure 
bismuth, a solution of its nitrate is precipitated with much water, 
which gives an insoluble basic nitrate, and this, when washed, 
dried, and reduced with a mixture of sodium carbonate and char¬ 
coal, yields the pure metal. 

Bismuth is a greyish-white metal, with a reddish tinge ; it is 
brittle, and can therefore be easily powdered, and crystallizes very 
easily. If bismuth is fused in a crucible, and when partly solidi¬ 
fied, the crust at the top broken, the still liquid metal poured out, 
and the crucible broken open when cold, its sides are found 
covered with beautiful iridescent crystals of bismuth, just as 
crystals of monoclinic sulphur are obtained in the same way. The 
colour of the crystals is due to a thin layer of oxide, with which the 
metal becomes covered when the air enters the crucible. The 
crystals are rhombohedra, the angle of which so closely approaches 
that of the cube, that bismuth was formerly thought to crystallize 
in the regular system. 

Bismuth has a specific gravity of 9*8, and melts at as low a tem¬ 
perature as 264°; it shares with water the property of expanding 
considerably on passing from the liquid to the solid state. At high 

L L 



5 H Texi-Book of Inorganic Chemistry. 

temperatures it volatilizes, but cannot be distilled, and burns when 
heated strongly in the air, producing yellow bismuth oxide. Hydro¬ 
chloric acid does not attack the metal, and sulphuric acid scarcely 
dissolves it, but it is at once dissolved by nitric acid, or by aqua 
regia, and converted into bismuth nitrate or chloride respectively. 

Bismuth unites with oxygen in three proportions and forms the 
compounds bismuth suboxide : BiO, bismuth oxide : Bi 2 O s , and 
bismuth pentoxide or bismuthic anhydride (unknown in the free 
state) : Bi 2 0 5 . Of these, bismuth oxide is the only important com¬ 
pound. 

Bismuth Suloxide : BiO, separates as a grey precipitate when 
caustic soda is added to a mixture of bismuth chloride and 
stannous chloride. By the action of the caustic soda, stannous 
hydrate and bismuth hydrate are first produced, and the former 
then abstracts oxygen from the latter, forming stannic hydrate, 
which remains dissolved in the excess of soda, and bismuth sub¬ 
oxide. Like other suboxides, it is converted by acids into a bismuth 
salt and free bismuth. 

Bismuth Oxide : Bi 2 0 3 , remains as a yellow powder, similar to 
lead oxide, when basic bismuth nitrate is heated to redness. It 
melts at a higher temperature, and then solidifies to a crystalline 
mass on cooling. It dissolves easily in hydrochloric, nitric, or sul¬ 
phuric acid, producing the corresponding bismuth salts. 

Bismuth Hydrate : BiO(OH), is thrown down as a white preci¬ 
pitate when a solution of a bismuth salt is mixed with caustic soda. 
It dissolves easily in acids, but not in an excess of caustic soda. 
When dry it forms a white powder, which is converted into yellow 
bismuth oxide on heating. 

Bismuthic Acid : Bi 2 0 5 + xH 2 0 , has not been much investigated, 
and is possibly a peroxide. It is obtained as a red heavy powder 
when chlorine is led through dilute caustic soda containing white 
bismuth hydrate in suspension, and is afterwards freed from alkali 
by washing with a dilute acid. It evolves oxygen when heated and 
sets free chlorine from hydrochloric acid. 

Bismuth oxide, like other sesquioxides, is only a weak base, and 
is partly separated from its salts by water, which in this case plays 


Bismuth. 


515 

the part of a base. Only a small number of bismuth salts are 
known. 

Bismuth Nitrate : (N 0 .,) 3 0 3 Bi + 5 H 2 0 .— Metallic bismuth dis¬ 
solves easily in nitric acid, with evolution of nitric peroxide, and on 
evaporating the clear solution large colourless crystals of bismuth 
nitrate of the above composition separate out. The salt dissolves 
in a small quantity of water, especially if a little nitric acid is added, 
but a large quantity decomposes the salt like antimony trichloride, 
and produces a white crystalline precipitate consisting of delicate 
silky scales. This substance is basic bismuth nitrate , of the com¬ 
position :— 

N 0 2 * OBi(OH)o, or (N 0 ,) 3 0 3 Bi + 2Bi(OH) 3 . 

It is largely used in medicine in cases of cholera and chronic 
diarrhoea, and is also employed as a cosmetic. 

Bismuth Chromate is deposited as a bright yellow precipitate, 
resembling lead chromate, when a solution of bismuth nitrate is 
mixed with one of potassium chromate. It is distinguished from 
the lead salt by its insolubility in a large excess of caustic soda. 

Bismuth Chloride : BiCl 3 .—If bismuth is heated in a retort 
and dry chlorine led into it, it burns and forms bismuth chloride, 
which distils over as a viscid liquid, and solidifies on cooling to a 
white crystalline mass. It boils at about 430°. The same com¬ 
pound may also be obtained by dissolving bismuth in aqua regia, 
evaporating to dryness, and distilling. It deliquesces in the air, 
and is decomposed by water, like the nitrate, producing a white 
crystalline powder of insoluble basic bismuth chloride (bismuth 
oxychloride') : BiOCl. 

Bismuth Sulphide : Bi 2 S 3 , occurs in nature as the mineral 
bismuthinite, and is thrown down as a dark-brown precipitate 
when sulphuretted hydrogen is led into a solution of a bismuth 
salt. It may also be prepared by fusing together bismuth and 
sulphur. 

Detection of Bismuth Compounds.— Bismuth, like lead, may 
be easily reduced by heating any of its compounds with sodium 
carbonate on charcoal in the reducing flame of the blowpipe, and 
the globules of metal are easily oxidized in the outer flame to yellow 

l 1 . 2 


516 


Text-Book of Inorganic Chemistry 


bismuth oxide, resembling lead oxide. The globules of reduced 
bismuth are, however, brittle, while those of lead are soft and 

''^Sdutions of bismuth salts give, like lead, a black precipitate with 
sulphuretted hydrogen, but a solution of a bismuth salt gives a 
white precipitate of a basic salt when poured into water, and yields 
no precipitate when mixed with sulphuric or hydrochloric acid. 
Further, the precipitate of bismuth hydrate obtained with caustic 
soda is not soluble in excess. 

Bismuth oxychloride may be distinguished from the corres¬ 
ponding compound of antimony by its insolubility in a strong 
solution of tartaric acid, in which the latter easily dissolves. 


TIN. 

Chemical Symbol : Sn .—Atomic Weight'. 118. 

Tin belongs to those few metals which have been known from 
the earliest times, and this is all the more remarkable as the metal 
never occurs native, and its ores are only sparingly found in nature. 
For the extraction of tin only one mineral is used—viz. cassitemte 
or tinstone , consisting of stannic oxide, with the composition : 
SnO,. This mineral is found in large quantities in Cornwall, in the 
Erzgebirge (Saxony and Bohemia), and other European localities, 
as well as in Banca, the Malay Peninsula, Peru, and Australia. 
The extraction of the metal from its ore is a very simple process. 
The tin-stone is first roasted to get rid of arsenic and sulphur, then 
stamped to powder and washed with water to remove the lighter 
earthy impurities, and finally smelted with anthracite or powdered 
charcoal in a reverberatory furnace. The reduced tin, u'hich sinks 
to the bottom of the hearth, is then withdrawn and refined by re¬ 
melting in iron pots and stirring with a green pole. Commercial 
tin is always more or less impure, and often contains arsenic. The 
pure metal is easily obtained by heating pure stannic oxide and 
charcoal in a crucible, with the addition of a little borax or some 
similar flux. 

Tin is a white metal, resembling silver, with a brilliant lustre, 
and a specific gravity of 7*3. It is soft and very malleable, so that 
it can be hammered out into thin foil (tin-foil), but possesses little 
tenacity. Up to ioo° its malleability increases, but w'hen heated 



Tin. 


517 


above this it becomes less malleable, and is even brittle at 200°. 
Advantage is taken of this property to produce what is called 
grain-tin — i.e. tin which has been heated to about 200° and broken 
to pieces under the hammer. Tin melts at 228°, and when broken 
shows a crystalline fracture. It is well known that a bar of tin 
when bent emits a peculiar noise—the so-called cry of tin , which 
is probably produced by the friction of the crystals of tin over one 
another. Commercial tin, when grasped in the hand, imparts to the 
skin a peculiar odour, not possessed by the metal itself. 

At the ordinary temperature, tin remains unchanged in the 
air, but becomes covered with a thin crust of white oxide when 
exposed to the air in the fused state. At a white heat it burns 
with a bright white light. Hot concentrated hydrochloric acid 
dissolves tin with evolution of hydrogen and produces stannous 
chloride ; the metal also dissolves in hot strong sulphuric acid, 
forming stannous sulphate and giving off sulphurous acid. Aqua 
regia dissolves the metal and forms stannic chloride. Ordinary 
concentrated nitric acid oxidizes the metal, with the production 
of large quantities of brown fumes and formation of white stannic 
oxide, insoluble both in water and nitric acid. Finally, tin is 
also dissolved when heated with caustic potash; hydrogen is 
then set free and a solution of potassium stannate produced. 

The compounds of tin are divided into two classes—the stan¬ 
nous and the stannic compounds. In the former the metal is a 
dyad, in the latter a tetrad. 

Stannous Compounds. 

Stannous Oxide : SnO.—The precipitate of white stannous 
hydrate obtained when caustic soda is added to stannous chloride 
easily dissolves in an excess of soda, and produces sodium stannite. 
If the alkaline solution is boiled it undergoes decomposition, and 
a black crystalline precipitate of stannous oxide is thrown down. 
After washing and drying, the compound remains unchanged in 
the air. If, however, the solution of stannous chloride is precipi¬ 
tated with sodium carbonate instead of caustic soda, the white 
stannous hydrate : Sn(OH) 2 , which is now produced (together with 
free carbonic acid) behaves differently and oxidizes in the air to 
stannic oxide. 

The oxysalts of stannous oxide are but little known. Stannous 
sulphate : S 0 2 * 0 2 Sn, may be prepared by dissolving tin in hot con- 


518 Text-Book of Inorganic Chemistry. 

centrated sulphuric acid, and is then deposited in crystalline plates 
when the solution cools. When its aqueous solution is boiled it 
undergoes decomposition, with formation of a basic salt. 

Stannous Chloride : SnCl 2 .—Tin when heated in dry hydro¬ 
chloric acid gas is gradually converted into a white crystalline mass 
of stannous chloride, which melts at 250° and sublimes at a higher 
temperature. Stannous chloride is, however, usually prepared by 
dissolving tin in hydrochloric acid, an operation which may be con¬ 
siderably accelerated by the addition of a few scraps of platinum 
foil or of a few drops of platinum chloride, which is at once reduced 
to metallic platinum. The contact of the platinum with the tin 
makes the latter far more electro-positive and causes it to dissolve 
more quickly. 1 On evaporating the solution of stannous chloride 
the salt crystallizes out with two molecules of water as the com¬ 
pound : SnCl 2 + 2 H 2 0 . This salt is manufactured on a large scale, 
and is brought into trade under the name of tin-salt. The readi¬ 
ness with which stannous chloride passes into stannic chloride 
makes it a powerful reducing agent. Mercuric chloride is at once 
reduced by it to mercurous chloride, and on warming to metallic 
mercury, while the stannous chloride becomes stannic chloride. 
An aqueous solution of stannous chloride becomes turbid when ex¬ 
posed to the air, and gradually deposits a basic stannic chloride. 

Stannous chloride is largely used in dyeing and for purposes of 
reduction. 

Stannous Sulphide : SnS.— On heating a mixture of tin-filings 
with an excess of sulphur in a closed crucible the two substances 
unite and form dark-grey, lustrous scales of stannous sulphide. 
The same compound is obtained as a dark-brown precipitate when 
sulphuretted hydrogen is led into a solution of stannous chloride. 
The precipitate dissolves in hydrochloric acid with evolution of 

1 The action of the platinum, which is deposited in a finely divided state on 
the tin, may be thus explained. Each particle of the platinum forms a galvanic 
couple with a corresponding particle of tin, and a series of tiny electric currents 
circulate between the particles of the two metals and the liquid in their neigh¬ 
bourhood. The production of these currents necessitates a more rapid solution 
of the tin than when the platinum is absent, and they are consequently not pro¬ 
duced. Other similar couples , as this arrangement has been called, are also 
occasionally used. Thus, a copper-zinc couple, prepared by treating zinc foil 
with a very dilute solution of copper sulphate, reduces many substances which 
zinc alone does not attack, or only very slowly.— Ed. 


Stannic Compounds ; 519 

sulphuretted hydrogen, but is insoluble in pure, freshly-prepared 
ammonium sulphide, and is so distinguished from stannic sulphide, 
which at once gives a soluble compound with ammonium sulphide. 
Stannous sulphide at once dissolves in ordinary yellow ammonium 
sulphide, because this compound contains dissolved sulphur, which 
converts the stannous sulphide into stannic sulphide, and this dis¬ 
solves in the ammonium sulphide. 

Stannic Compounds. 

Stannic Oxide : Sn 0 2 , constitutes the mineral cassiterite or tin¬ 
stone , and is found either in compact masses or else in brown quad¬ 
ratic' crystals. It may be obtained as a white amorphous powder 
by burning tin in the air or by heating stannic hydrate. In the 
crystalline state it is formed when the vapours of stannic chloride 
and water are led through a red-hot porcelain tube. Stannic oxide 
is insoluble in acids, and is not attacked when fused with acid 
potassium sulphate, but may be converted into soluble sodium 
stannate by fusion with caustic soda or sodium carbonate. Stannic 
oxide has far more the character of an acid than of a base. 

Stannic Hydrate. —Several hydrates of stannic oxide are known, 
and are all weak acids and very weak bases. If tin is warmed 
with moderately strong nitric acid, a violent reaction ensues, tor¬ 
rents of nitric peroxide are evolved, and a white insoluble powder 
of inetastannic acid ', of the probable composition : Sn(OH) 4 , re¬ 
mains behind. This hydrate is insoluble in water and nitric acid, 
but is converted by strong hydrochloric acid into a gelatinous mass 
soluble in water. The solution in water contains stannic chloride, 
and on the addition of sulphuric acid a white precipitate of stannic 
sulphate is obtained. This is an unstable compound, from which 
water again abstracts all the sulphuric acid. Metastannic acid, 
when heated to ioo°, loses water and becomes converted into 
a compound having the composition : SnO(OH) 2 , but which is 
different from stannic acid ; at a red heat both this compound 
and metastannic acid are converted into stannic oxide. Meta¬ 
stannic acid when digested with cold caustic soda forms sodium 
metastannate , a slightly soluble salt of which the composition is 
uncertain. But if the hydrate is boiled with caustic soda an 
easily soluble salt is formed, which can be obtained in the 
crystalline form, and which is known as sodium stannate , of the 
composition : SnO(ONa) 2 . The same salt is also produced when 


520 Text-Book of Inorganic Chemistry\ 

stannic oxide is fused with caustic soda. It is largely used in 
calico-printing. 

The corresponding potassium stannate : SnO(OK) 2 + 3 H 2 0 , is 
obtained in the same way as the sodium compound. It is insoluble 
in caustic potash, but soluble in water, and is deposited in colourless 
crystals ; its solution reacts strongly alkaline. 

A second variety of stannic hydrate is known under the name 
of stannic acid , the composition of which is probably : SnO(OH) 2 , 
corresponding to the above stannates. It is obtained as a white, 
voluminous precipitate when stannic chloride is mixed with sodium 
carbonate. Stannic acid is insoluble in water, but dissolves easily 
in strong hydrochloric or nitric acid as well as in caustic soda. 
The acid solution when boiled gradually deposits stannic acid, 
provided too much free acid is not present. Stannic acid, when 
kept for some time under water, becomes insoluble in nitric acid. 

Stannic Chloride : SnCl 4 .—This volatile liquid, which was 
previously known as Spiritus Libavii finnans , is obtained when 
tin is heated in a retort and chlorine passed over it. The dis¬ 
tillate when purified by redistillation is a colourless heavy liquid 
of specific gravity 2-28, and boiling at 114° It fumes strongly in 
moist air, almost as much as sulphuric anhydride, and hisses when 
dropped into water, forming the hydrated salt: SnCl 4 + 3H.,0, 
which crystallizes out. A dilute aqueous solution is decomposed 
on boiling, with separation of stannic hydrate. A solution of stannic 
chloride is best obtained by dissolving tin in aqua regia. 

Stannic chloride unites with the alkaline chlorides and forms 
crystalline double salts. The double compound with ammonium 
chloride: SnCl 4 + 2 NH 4 C 1 , is used by calico-printers under the 
name of pink-salt. 

Stannic Sulphide : SnS 2 , is obtained as a yellow amorphous 
precipitate when sulphuretted hydrogen is led through a not too 
acid solution of stannic chloride. It is insoluble in water, but 
dissolves in strong hydrochloric acid. Hot concentrated nitric 
acid oxidizes it to stannic hydrate and sulphuric acid. Stannic 
sulphide possesses the character of a sulpho-acid, and as such easily 
dissolves in ammonium sulphide, producing soluble ammonium 
sulphostannate , which is easily decomposed by acids with re-sepa¬ 
ration of stannic sulphide. 

Stannic sulphide is prepared in the dry way as golden scales 


Stannic Compounds. 5 21 

by heating together tin and sulphur with volatile substances such 
as mercury and ammonium chloride—which, by their volatilization, 
appear to prevent the temperature rising high enough to decompose 
stannic sulphide into stannous sulphide and sulphur. A solid 
amalgam of tin is first prepared by adding six parts of mercury to 
twelve parts of molten tin. This is powdered when cold, mixed 
with seven parts of flowers of sulphur and six parts of ammonium 
chloride, and the mixture introduced into a glass flask, the mouth 
of w r hich is loosely closed by piece of chalk. The flask is next 
placed in a crucible surrounded with sand and heated in a good 
draught to low redness. During the process mercuric chloride, 
ammonium and stannic chlorides sublime, and at the bottom 
of the flask is found the stannic sulphide as a loose mass of a 
golden-yellow colour. The substance so prepared is called mosaic 
gold; it is much more stable than the precipitated sulphide, being 
scarcely attacked by strong hydrochloric or nitric acid. 

Mosaic gold is used for a number of technical purposes, but 
chiefly for cheap gilding. 

More important than the technical applications of the com¬ 
pounds of tin are those of the metal itself, either alone or alloyed 
with other metals. Considerable quantities of tin are consumed in 
the manufacture of tin-plate—i.e. iron plates coated with a thin 
layer of tin ; copper vessels are tinned internally, and tm-foil is 
largely used for packing small articles. Tin-foil usually contains 
traces of copper, which gives it greater tenacity. Not unfrequently 
it also contains lead, which may produce very deleterious effects if 
articles of food are packed in the foil. Not only is lead more 
easily attacked by the acids in the food than tin, but the compounds 
of lead are far more injurious when introduced into the system than 
those of tin. 

Alloys of Tin.— Different alloys of lead and tin-known as 

■bewter _are used for the manufacture of various vessels, and as a 

basis for the cheaper silver-plated articles ; common pewter con¬ 
tains about four parts of tin to one of lead. Common soft-solder, 
chiefly used for soldering tin-plate, is an alloy of 2 parts of tin 
with i part of lead, and melts at about i8o°, or of equal parts of the 
two metals, with a melting-point of about 200°. Britanma-metal , 
a better kind of pewter, is essentially an alloy of about 9 parts 
of tin with 1 part of antimony. The various alloys of tin with 


522 Text-Book of Inorganic Chemistry. 

copper, bronze, &c., will be described under the latter metal; 
an amalgam of tin is used for silvering looking-glasses. By melt¬ 
ing together tin with other easily fusible metals—<?.£•. lead, bis¬ 
muth, and cadmium—alloys are obtained with very low melting 
points, and which are largely used in the process of stereotyping. 
An alloy of i part of lead, i part of tin, and 2 parts of bismuth 
(.fusible metal ), melts at 95 0 ; and one consisting of 4 parts of 
bismuth, 2 parts of lead, 1 part of tin, and 1 part of cadmium 
( Wood's metal), melts at about 68°. 

Detection of Tin Compounds. 

Tin and antimony are the only elements which are converted 
by nitric acid into insoluble white oxides \ of these the antimony 
compound may be distinguished by its solubility in a solution of 
tartaric acid, which leaves the stannic oxide undissolved. Tin itself 
may be distinguished from antimony by its solubility in hydro¬ 
chloric acid, which does not attack the latter substance. Stannous 
compounds, in consequence of the readiness with which they pass 
into the stannic form, are powerful reducing agents. Stannous 
chloride precipitates mercury and gold from solutions of their salts, 
the latter even in the cold. With sulphuretted hydrogen stannous 
salts give a dark-brown precipitate of stannous sulphide, and the 
same reagent throws down yellow stannic sulphide when led 
through solutions of the stannic salts. Both sulphides dissolve in 
yellow ammonium sulphide. Stannic oxide has more the character 
of an acid than a base ; even stannous oxide possesses faint acid 
properties, and unites both with strong acids and with strong bases 
to form salts. Metallic tin, as white malleable globules, may be 
reduced from its compounds when a compound of the metal is 
mixed with sodium carbonate and heated on charcoal in the inner 
blowpipe flame, though the reduction is rather more difficult than 
that of lead. At the same time, an indistinct white incrustation is 
formed on the charcoal. Compounds of tin do not tinge a bead of 
borax, but if a minute trace of copper is added, and the bead held 
in the reducing flame of the blowpipe, it easily becomes of a fine 
red colour, due to reduced cuprous oxide. 



Copper . 


523 


COPPER. 

Chemical Symbol : Cu.— Atomic Weight'. 63'4. 

Copper, which has been known from the earliest times, is one 
of the few metals which possess a decided tint ; it is of a bright red 
colour. The metal is sometimes found in the free state in nature, 
often well crystallized in octahedra, but more commonly occurs in 
combination with other elements. The most important ores of 
copper are : cuprite or red copper ore (cuprous oxide : Cu 2 0 ), 
copper-glance (cuprous sulphide : Cu 2 S), copper-pyrites (sulphide of 
copper and iron : CuFeS 2 ), purple copper ore of similar composition, 
and tetrahedrite or fahlerz (a mixed sulphide of copper, antimony, 
arsenic, &c.). Besides these, some oxy-salts of copper are also 
found in the mineral kingdom, among which may be mentioned 
the basic carbonates : malachite and azunte. 

The extraction of copper from its oxygen compounds (cuprite, 
malachite, &c.) is very simple, and analogous to the method used 
for the reduction of iron. The roasted ores are heated in a blast 
furnace with some form of carbon, and the reduced copper collects 
on the hearth under the slag. Only small quantities of copper are 
extracted from these comparatively rare ores, far more is obtained 
from copper-pyrites, a double sulphide of copper and iron. The 
method employed for the separation of copper from these ores is a 
complex one ; a whole series cf processes is necessary to com¬ 
pletely separate both the iron and the sulphur. These processes, 
which depend upon the attraction of copper for sulphur and iron 
for oxygen, consist in alternately roasting and melting the orea. 
The roasting gradually removes the sulphur and oxidizes the iron, 
which unites with silica to form a fusible slag. From the copper 
sulphide which then remains, metallic copper is obtained by oxi¬ 
dizing the.sulphur. 

Commercial copper is never quite pure, although the impurities 
| are usually insignificant. The pure metal may be obtained as a 
red powder by heating pure copper oxide in a stream of hydrogen. 

Copper melts* only at a high temperature, above i,ooo°, and 
burns when heated in the oxy-hydrogen flame. Its specific gravity 
is about 8-9. The metal takes a good polish, is malleable and 
I ductile, and very tenacious. Bright copper remains unchanged in 
dry air, but in the presence of moisture becomes covered with a 
green layer of basic copper carbonate. This green rust of copper 




524 Text-Book of Inorganic Chemistry . 

is often called verdigris , which is really a basic acetate of the metal. 
Copper is more rapidly corroded if the surfaceds moistened with a 
dilute acid and exposed to the air; it is also attacked by ammonia 
or a solution of common salt in the presence of air. In the absence 
of air neither hydrochloric nor dilute sulphuric acid has any action 
on the metal. Hydrochloric acid does not dissolve it even on 
heating, but hot concentrated sulphuric acid oxidizes it to copper 
sulphate, and sulphurous anhydride is set free (p. 151). Unlike iron, 
zinc, and other metals, copper cannot decompose water and form 
copper oxide and hydrogen ; even when copper is heated to redness 
in steam no change is produced. If copper is heated to redness in 
the air it becomes covered with a layer of black cupric oxide which 
separates as scales when the metal is hammered. This copper 
scale consists chiefly of cupric oxide, but contains some metallic 
copper and some cuprous oxide. 

The best solvent for copper is nitric acid, which dissolves it with 
evolution of nitric oxide (p. 193), and produces a blue solution of 
cupric nitrate. Copper is deposited from its solutions by more 
positive metals— e.g. zinc or iron—and also separates in a compact 
form at the negative pole when an electric current is passed through 
a solution of the sulphate. The process of producing electrotypes 
or copies of various objects in copper depends upon this deposition 
of the metal by an electric current. 

Compounds of Copper. 

Copper forms with oxygen, sulphur, and the halogens, two series 
of compounds ; the one series contains a single atom of dyad 
copper, and the other 2 atoms of copper which also play the part 
of a dyad. The two series of compounds are called cupric and 
cuprous compounds respectively—^, cupric chloride : CuCl 2 , and 
cuprous chloride : Cu 2 Cl 2 . Of these two series, the cupric com¬ 
pounds are the most important, and shall therefore be first dis¬ 
cussed. All compounds of copper are poisonous, especially those 
which are soluble in water. 

Cupric Compounds. 

Cupric Oxide, Black Oxide of Copper : CuO, is a black amor¬ 
phous powder, obtained by glowing finely divided copper in the air, 
or better by heating cupric nitrate. Copper oxide is largely used 
in organic analysis to oxidize organic substances (p. 289). For this 
purpose it does not require to be pure, but should be in the form of 


Cupric Compounds. 5 2 5 

scales, and is best prepared from copper scale. A layer of the scale, 
which has been freed from mechanical impurities, is moistened 
with nitric acid (containing no chlorine) in a Hessian crucible, then 
a second layer added and also moistened with the acid until the 
crucible is about three parts full. The whole is then heated to low 
redness until all the nitric acid is destroyed and expelled. The 
nitric peroxide set free during the decomposition of the cupric 
nitrate suffices to oxidize the copper and cuprous oxide, contained 
in the scale, to cupric oxide, and if the temperature has not been 
raised high enough to cause the oxide to cake together, it is ob¬ 
tained in the form of scales well suited for organic analysis. 

Cupric oxide may be obtained in the wet way as a black pre¬ 
cipitate, by adding hot caustic soda to a hot solution of a cupric 
salt. Cupric oxide is completely insoluble in water, but readily 
dissolves in hydrochloric, nitric, or sulphuric acid, without any 
evolution of gas, then forming the corresponding cupric compounds. 
When heated in a stream of hydrogen it easily parts with its oxygen 
and yields metallic copper as a red powder. 

Cupric Hydrate : Cu(OH).,.—If a solution of a cupiic salt is 
mixed with an excess of caustic soda at the ordinary temperature, 
a blue voluminous precipitate of copper hydrate is obtained, in 
which the copper oxide and water are so loosely combined that the 
compound is decomposed at ioo°, and black cupric oxide even 
separates slowly at the ordinary temperature. If insufficient caustic 
soda is employed a basic salt is obtained, which is also of a blue 
colour, but which remains unchanged when boiled. 

Cupric Sulphate, Blue Vitriol'. S 0 . 2 - 0 . 2 Cu + 5 H., 0 , is de¬ 
posited, when its aqueous solution is evaporated, as large blue 
crystals belonging to the triclinic system, and is the salt from 
which nearly all the other compounds of copper are obtained. 
Copper sulphate may be prepared by dissolving copper oxide in 
sulphuric acid, or by heating copper with concentrated sulphuric 
acid. Considerable quantities of copper sulphate are produced by 
the oxidation of copper-pyrites when roasted in the air, but always 
mixed with morq or less ferrous sulphate. It cannot be completely 
freed from this impurity by simple recrystallization as the two salts 
crystallize together, but if sufficient nitric acid is added to oxidize 
the iron to the ferric state, the solution evaporated to dryness, and 
the dry mass extracted with water, most of the iron remains behind 
as a basic salt. Copper is precipitated from the blue solution of 




526 Text-Book of Inorganic Chemistry . 

cupric sulphate, produced in mines where copper-pyrites is worked, 
by adding scrap iron. The metal so obtained is known as cement- 
copper. 

Cupric sulphate is easily soluble in water ; one part of the salt 
dissolves in 2-5 parts of cold water or in 0-5 part of boiling water. 
The blue crystals slightly effloresce when exposed to the air, and 
lose 4 molecules of their water of crystallization at ioo°, but the 
fifth molecule of water (water of constitution) is only expelled at 
about 200 0 . The anhydrous salt is white, but regains its blue colour 
in contact with water, being reconverted into the hydrated com¬ 
pound. The energy with which anhydrous cupric sulphate unites 
with water is so great that we employ it to abstract the last traces 
of water from 99 per cent, alcohol, and to remove traces of water 
from many other liquids in which the salt does not dissolve. 

Cupric sulphate unites with the sulphates of the alkalies and 
forms double salts, which may easily be obtained in the crystalline 
form combined with 4 molecules of water. They are of a lighter 
colour than cupric sulphate and less soluble in water. 


Cupric Ammonium Sulphates.— If a concentrated aqueous solu¬ 
tion of cupric sulphate is mixed with strong ammonia until the 
precipitate of cupric hydrate which is first produced is redissolved, 
a deep blue liquid is produced, which, on standing, or better, when 
covered with a layer of alcohol, deposits a layer of deep blue crys¬ 
tals of the composition: S 0 2 - 0 2 Cu + 4NH3 + H. 2 0. These crys¬ 
tals gradually decompose and become opaque when exposed to the 
air, and when heated to 150° lose their water with 2 molecules of 
ammonia, leaving a green powder of the composition : SO,- 0 . Cu + 

2NH3. . ~ ' 

The formation and constitution of these two compounds may 
be represented as follows (see p. 563). On the addition of ammonia 
to cupric sulphate, ammonium sulphate and cupric hydrate are first 
produced, and the latter compound then appears to unite with 
ammonia and form cupri-diammonium hydrate :— 


Cu 


OH 

OH 


+ 2NH, = 


OH-NH. 

OH-NE 


Cu, 


or the hydrate of a hypothetical diammonium in which 2 atoms 
of hydrogen are displaced by 1 atom of dyad copper. This sub¬ 
stance does not appear to exist in the free state, but unites at 
once with the ammonium sulphate, and produces the blue salt, 
diammon-cupri-diammonium sulphate , thus :— 


Cupric Compounds. 


527 


SO ! 0NH 4 . OH-NH s ) c fONHg-NHg) ^ oH n 

bU 2 (ONH 4 + OH-NH ? f ^ u “ ( ONH 3 - NH3I Cu + 2H ~°’ 


Ammonium 

sulphate 


Cupri-diammonium 
hydrate 


Diammon-cupri- 

diammonium 

sulphate 


which may be considered as ammonium sulphate in which two 
atoms of hydrogen have been displaced by the dyad radical cupri- 

diammonium,'. nh 3 }cu- On heating the blue compound to 150° 

it parts with two molecules of ammonia, and leaves the green 
powder mentioned above, which is cupri-diammonium sulphate :— 

f ONH..-NHJ r crk ( ONHJ r -iwt u 

S ° 2 | ONHg-NH^ [ Cu = S ° 2 ] ONH3 1 Cu + 2NH2 - 


Diammon cupri- 
diammonium 
sulphate 


Cupri-d ammonium 
sulphate 


The constitution of the ammoniacal compounds of cobalt 
(p. 487) is possibly similar to that of the above salts. 


Cupric Nitrate : (N 0 2 ) 2 0 2 Cu + 3H 2 0, is easily obtained by 
dissolving copper in nitric acid, and separates out when the solu¬ 
tion is concentrated as dark blue deliquescent crystals, easily 
soluble both in water and in alcohol. 


Cupric Phosphate : (PO) 2 0 6 Cu 3 + 3H0O, is thrown down as a 
blue-green precipitate when ordinary sodium phosphate is added 
to a solution of cupric sulphate. Various phosphates of copper 
occur in the mineral kingdom. 


Cupric Arsenite, probably: As-J 


0 .,Cu 

OH 


is obtained as a green 


precipitate varying in tint when a solution of sodium arsenite is 
poured into one of copper sulphate, with constant stirring. This 
compound is used as a pigment under the name of Scheelds green. 
A second green pigment—known as Schweinfurt-green— is a double 
compound of copper arsenate and acetate. Both these compounds 
are highly poisonous, and their use is therefore very limited. 


Cupric Carbonates.— All the known carbonates of copper are 
basic compounds, even when normal copper sulphate is mixed with 
normal sodium carbonate, the bluish-green gelatinous precipitate, 
which after a time changes to a bright green powder, consists of 
a basic carbonate of the composition: C 0 0 2 Cu + Cu(OH) 2 = 


528 


Text-Book of Inorganic Chemistry. 


CO | ocuOH Tlie same com P ound * s f° und nature > espiceally 
in Siberia, as the beautiful mineral, malachite, which in its massive 
form is used for the manufacture of vases, tops of tables, &c. 

A second basic carbonate of copper is the dark blue mineral 
azurite , of the composition : 2CO - 0 2 Cu + Cu(OH) 2 . 

Cupric Acetates .—These compounds belong strictly to organic 
chemistry, but it may be remarked that normal cupric acetate is 
soluble in water and is deposited from its solution in blue crystals. 
The basic acetates are insoluble or difficultly soluble in water. One 
of these, or rather a mixture of several, is the poisonous substance 
called verdigris. 

Cupric Chloride : CuCl 2 + 2H a O,is obtained by dissolving cupric 
oxide in hydrochloric acid, and on concentrating the solution is 
deposited in small, green, rhombic prisms. When heated the salt 
loses its water and the anhydrous chloride remains behind as a 
yellowish-brown powder, which rapidly attracts moisture from the 
air and deliquesces. When more strongly heated, it is decomposed 
into cuprous chloride and free chlorine. 

Cupric chloride unites with cupric hydrate and forms basic 
compounds. One of these is the mineral atacamite, and has the 
composition: CuCl 2 + 3Cu(OH) 2 . These insoluble basic chlorides 
may be obtained by mixing a solution of cupric chloride with a 
quantity of caustic soda insufficient to convert all the copper into 
cupric hydrate. 

Cupric Bromide : CuBr 2 , closely resembles the chloride. Cupric 
Iodide, strange to say, has not been prepared. If potassium iodide 
is added to a solution of cupric sulphate, a white precipitate of 
cuprous iodide is produced and iodine is set free. 

Cupric Sulphide : CuS, is thrown down as a black precipitate 
when sulphuretted hydrogen is led through a solution of a cupric 
salt. In the moist state it readily attracts oxygen from the air and 
is converted into cupric sulphate, and must therefore be washed 
with water containing sulphuretted hydrogen and dried in a vacuum. 
When heated in a stream of hydrogen it parts with one-half of its 
sulphur and is converted into cuprous sulphide : Cu 2 S. Cupric 


Cuprous Compounds. 529 

sulphide is found in nature as the mineral covellite in dark-blue 
hexagonal prisms. 


Cuprous Compounds. 

Cuprous Oxide : Cu., 0 .—This compound is found in nature 
as the mineral cupj'ite or red copper ore either in compact crys¬ 
talline masses, or else in small regular octahedra. The crystalline 
compound may be artificially prepared by fusing solid cuprous 
chloride with dry sodium carbonate. It then remains as a red 
crystalline powder when the fused mass is afterwards extracted 
with water. It may also be prepared in the same form but in 
smaller crystals, by gently warming a solution of cupric sulphate 
with an excess of caustic soda and grape sugar. The grape sugar 
is then oxidized by half the oxygen contained in the copper oxide. 
Cuprous oxide remains unchanged in the air and is quite insoluble 
in water. When treated with hydrochloric acid, it is converted 
into white cuprous chloride, which dissolves in excess of the acid, 
and forms a colourless solution in the absence of air. Dilute 
sulphuric acid converts it into cupric sulphate with the separation 
of one-half of the copper as a red powder. Other oxy-acids behave 
in the same way, except nitric acid, which completely dissolves it. 

Cuprous Hydrate'. Cu 2 (OH) 2 .—When a solution of cuprous 
chloride in hydrochloric acid is poured into an excess of caustic 
soda, cuprous hydrate is thrown down as a yellow crystalline 
precipitate. The compound is very unstable, rapidly absorbing 
oxygen from the air, and producing blue cupric hydrate. 

The cuprous oxy-salts are very unstable, and are scarcely known. 
The corresponding compounds with the halogens are, however, 
more permanent. 

Cuprous Chloride : Cu.,Cl. 2 .—This substance is a white crys¬ 
talline powder, insoluble in water, but soluble in hydrochloric acid, 
from which it crystallizes in white tetrahedra. Cuprous chloride 
may be obtained by boiling cupric chloride with excess of hydro¬ 
chloric acid and metallic copper in a stream of carbonic acid to 
exclude the air, the oxygen of which would oxidize the cuprous 
chloride thus produced. On afterwards pouring the clear acid 
solution into a large quantity of cold boiled water, the cuprous 
chloride is deposited as a white crystalline powder. Cuprous chlo- 

M M 


530 Text-Book of Inorganic Chemistry. 

ride may also be prepared by adding a solution of stannous chloride 
containing free hydrochloric acid to a solution of cupric chloride. 
It then crystallizes out also as white tetrahedra, while the stannic 
chloride, produced at the same time, remains dissolved in the acid 
liquid. 

Cuprous chloride when exposed to the air in the moist state 
becomes of a green colour, and at ioo° large quantities of oxygen 
are absorbed, forming the oxy-chloride : CuCl,,CuO. But if this 
compound is further heated (up to about 400°), it again parts with 
its oxygen and leaves cuprous chloride. By this means, consider¬ 
able quantities of oxygen may be obtained from the air. 

Ammonia, as well as hydrochloric acid, dissolves considerable 
quantities of cuprous chloride, and both solutions absorb carbonic 
oxide. The compound which is produced is deposited in colour¬ 
less scales when a strong solution of cuprous chloride in hydrochloric 
acid is saturated with the gas. The ammoniacal solution also 
absorbs and unites with certain gaseous hydrocarbons, especially 
acetylene. 

Cuprous Iodide : Cu 2 I 2 , is a white powder insoluble in water 
and dilute acids, and is deposited when potassium iodide is added 
to a solution of cupric sulphate containing sulphurous acid or 
ferrous sulphate, these two substances acting as reducing agents. 
The insolubility of cuprous iodide makes this salt useful for sepa¬ 
rating iodine from solutions of its salts. Unlike cuprous chloride, 
the iodide remains unaltered when exposed to the air. 

Cuprous Sulphide : Cu. 2 S, is obtained by heating finely-divided 
copper with sulphur, or by heating cupric sulphide in a stream of 
hydrogen. It forms a dark-grey mass, which melts at a red-heat, 
and when slowly cooled may be obtained in a crystalline form. 
Cuprous sulphide is found in nature as copper-glance , which crys¬ 
tallizes in fine rhombic prisms. It also occurs united with other 
sulphides— eg. with iron sulphide in copper-pyrites and purple- 
copper ore , as well as with the sulphides of silver, antimony, arsenic, 
&c., in other minerals. 

Alloys of Copper.— Few metals are so largely used as copper, 
both in the form of its salts, and as alloys with other metals. The 
most important of these alloys are those with zinc, tin, nickel, and 
silver respectively. 

Brass is an alloy of copper and zinc obtained by melting the 


Alloys of Copper. 531 

two metals together, and is of a darker or lighter colour accordingly 
as it contains more or less copper. Ordinary brass contains about 
70 parts of copper to 30 parts of zinc. A larger proportion of 
zinc makes an alloy of a pale yellow colour, known as Muntz-metai , 
and largely used for sheathing wooden-ships. By increasing the 
percentage of copper and adding a little tin, a more malleable alloy 
is obtained, which can be beaten into thin sheets, and is called 
tombac or Dutch metal. 

Gun-metal brotize and bell-metal are all alloys of copper and 
tin, with about 80 to 90 per cent, of copper, to which a little lead 
or zinc is often added. Bronze contains less copper than gun- 
metal, and bell-metal less than bronze. Speculum-metal is a hard 
white alloy of about 1 part of tin and 2 parts of copper, to which 
some arsenic is not unfrequently added. It takes a high polish, 
and is used for the mirrors in reflecting telescopes. 

An alloy of copper and nickel is used for small coins in 
Germany, Belgium, and some other countries; in Switzerland the 
small coins are made of an alloy of copper, nickel, zinc, and silver. 
German-silver or nickel-silver is an alloy of copper, nickel, and 
zinc (p. 492). All the silver used in the arts consists of an 
alloy of copper and silver ; and, finally, an alloy of copper and 
aluminium, containing about 9 parts of the former metal to 1 part 
of the latter, forms aluminium-bronze or aluminium-gold, a tena¬ 
cious, bright yellow alloy, largely used in the manufacture of small 
articles. 


Detection of Copper Compounds. 

Copper belongs to those metals of which the sulphides are in¬ 
soluble in dilute acids, and are separated when a stream of sulphu¬ 
retted hydrogen is led through an acid solution of one of their salts. 
The cupric sulphide then obtained is black and resembles lead, 
bismuth, and some other sulphides. This precipitation of cupric 
sulphide from acid solutions serves to distinguish copper from 
nickel, which it otherwise resembles in some points. A character¬ 
istic test for a solution of a copper salt is the bluish precipitate 
produced on the addition of ammonia and the dark blue solution 
obtained when the ammonia is added in excess. A solution of a 
copper salt, even when extremely dilute, gives, on the addition of 
potassium ferrocyanide, a reddish brown precipitate of cupric ferro- 
cyanide : Cu 2 FeCy 6 ; only uranium compounds form a precipitate 
of this colour with potassium ferrocyanide. 


M xM 2 


532 Text-Book of Inorganic Chemistry. 

Small quantities of a compound of copper may be easily de¬ 
tected by heating with borax on a platinum wire before the blow¬ 
pipe. A green bead is then formed when heated in the outer flame, 
which becomes colourless or of a dirty red colour (from cuprous 
oxide) in the reducing flame. The red cuprous oxide may be more 
readily obtained by the addition of a trace of stannic oxide, or any 
compound of tin, and then heating in the reducing flame. 


533 


NOBLE METALS . 

This group includes mercury, silver, gold, platinum, and five 
other rare metals allied to platinum ; they have been called noble 
because they do not tarnish in the air at the ordinary temperature. 
They are further distinguished by the fact that their oxides are all 
decomposed when heated into the metal and free oxygen, with the 
exception of those of osmium and ruthenium, which volatilize un¬ 
changed. 


MERCURY. 

Chemical Symbol : Hg .—Atomic Weight : 200. 

This metal, with the exception of gallium, is the only one which 
is liquid at the ordinary temperature, and as it possesses a white 
colour like silver, is sometimes called quicksilver. It is only 
sparsely distributed in nature, but large quantities are found in the 
few localities where it occurs. Mercury is sometimes found free in 
the mineral kingdom or alloyed with silver or gold, but far more 
commonly in combination with sulphur as cinnabar : HgS. The 
mercury mines of Idria in Austria and of Almaden in Spain have 
been worked for centuries, and considerable quantities of the metal 
are also obtained from China. In late years large deposits of cin¬ 
nabar have been discovered in California. 

To extract mercury from cinnabar on a small scale the pow¬ 
dered ore is mixed with quicklime and heated in a retort, the neck 
of which dips under water. During the heating calcium sulphide 
and mercuric oxide are formed, and the latter compound then breaks 
up into mercury and oxygen. The mercury then condenses under 
the water, and the oxygen partly escapes, but principally serves to 
convert the calcium sulphide into calcium sulphate. 

On the large scale mercury is obtained from cinnabar by 
roasting the ore with excess of air in specially constiucted furnaces, 



534 Text-Book of Inorganic Chemistry. 

and passing the vapours of sulphurous anhydride and mercury 
which are thus produced through a series of chambers or through 
rows of earthenware vessels, where the mercury condenses. 

Commercial mercury, even when it has been distilled, is never 
pure, but contains mechanical impurities as well as traces of zinc, 
lead, and other less volatile metals dissolved in it. The presence 
of these impurities may be easily detected by allowing a drop of 
the metal to flow over the hand, when it runs off clear if pure, but 
leaves a dirty, dull tail on the skin when contaminated with other 
metals. Mechanical impurities may be separated by filtering the 
mercury through a filter of writing-paper with a pin-hole at the 
base. To separate the dissolved metals the mercury is placed in a 
flat dish and covered with a layer of ordinary nitric acid. The 
acid then oxidizes some of the mercury and produces mercurous 
nitrate, which then transfers part of its oxygen to the metals mixed 
with the mercury. After the whole has stood for several days, with 
frequent stirring, the acid liquid is poured off, the metal washed 
with water, then run through a separating-funnel and dried. This 
method separates most of the impurities contained in the metal, 
but chemically pure mercury can only be obtained by other and 
more tedious processes. 

Mercury has a specific gravity of 13*55; it solidifies at -40° 
to a crystallline malleable mass, with a specific gravity of 14*4. 
Although the metal only boils at 357 0 , it volatilizes even at the 
ordinary temperature of the air; if apiece of gold leaf is suspended 
in a jar containing a little mercury it soon becomes white, and 
amalgamated from the mercury with which it has combined. 

Mercury remains unchanged in the air at the ordinary temper¬ 
ature, but when heated to about 300° it combines with oxygen and 
becomes changed into red mercuric oxide, which coats the metal. 
It is completely insoluble in water and is not attacked by hydro¬ 
chloric or dilute sulphuric acid. Concentrated sulphuric acid is 
also without action upon it at the ordinary temperature, but when 
heated dissolves it, producing mercuric sulphate and sulphurous 
anhydride. The best solvent tor mercury is nitric acid, with which 
it forms either mercuric or mercurous nitrate according to different 
conditions. 

The vapours of metallic mercury and mercury compounds, 
especially those soluble in water, are all poisonous. Even the 
lower animals cannot exist in an atmosphere containing traces of 
mercury vapour. Entomological collections may be effectually 


Mercury. 535 

preserved from destruction by minute animals if a few drops of 
mercury are placed in the cases containing the specimens. 


Compounds of IVIercury. 

Mercury, like copper, forms two series of compounds—the one 
containing a single atom of the dyad metal, and the other con¬ 
taining two atoms of mercury, which also perform the functions of a 
dyad. These two series of compounds are the mercuric and mer¬ 
curous compounds respectively. Many are highly valued and 
powerful medicines. 


mercuric Compounds. 

mercuric Oxide ! HgO.—If mercury is heated in the air to 
about 300°, it gradually becomes oxidized and converted into a 
red powder of mercuric oxide (p. 10). This change goes on too 
slowly to prepare any quantity of the oxide, and it is preferable to 
obtain it from its nitrate. For this purpose dry mercuric nitrate is 
mixed with an equal weight of mercury, and the mixture gently 
heated as long as acid vapours are given off. The oxide then 
remains as a heavy, red, crystalline powder, of specific gravity 
11 -2, and which becomes of a paler colour when rubbed. 

Besides this crystalline form of mercuric oxide, a second modi¬ 
fication may be obtained as an amorphous yellow powder, by 
precipitating a solution of mercuric nitrate with caustic soda. 
Mercury, like the other noble metals, is precipitated from its solu¬ 
tions by the alkalies and other strong bases as the oxide and not 
as the hydrate. Mercuric hydrate is unknown. 

Mercuric oxide, as previously stated (p. 10), breaks up into 
mercury and oxygen when heated. 

mercuric Sulphate : S 0 2 - 0 2 Hg, is obtained by dissolving 
mercuric oxide or metallic mercury in sulphuric acid, in the latter 
case with evolution of sulphurous anhydride. It unites with a 
small quantity of water to form a colourless crystalline compound 
containing 1 molecule of water. With a large quantity of water 
it is decomposed into free sulphuric acid and a basic sulphate of 
the composition : S 0 2 - 0 2 Hg + 2HgO, which separates as a heavy 
lemon-coloured powder. This compound was formerly used in 
medicine under the name of turpeth mineral. Mercuric sulphate 


536 Text-Book of Inorga, 7 iic Chemistry. 

is manufactured on a large scale, and used for the preparation of 
the two chlorides of mercury. 

Mercuric Nitrate : (N 0 2 ) 2 0 2 Hg + 8H..O, may be prepared by 
acting on mercury with an excess of nitric add, until a drop of the 
solution no longer gives a precipitate with hydrochloric acid—a 
proof that it contains no mercurous nitrate. The acid solution, 
when sufficiently concentrated, and when cooled to — 15 0 , deposits 
large, colourless, rhombic crystals of the above composition. This 
salt is unstable ; it melts in its water of crystallization at 7 0 , and 
changes into a basic salt of the composition : (N 0 2 ) 2 0 2 Hg + HgO + 
2H 2 0, which is also deposited when a hot concentrated solution of 
mercuric nitrate is allowed to cool. Both these salts are decomposed 
by water, yielding a still more basic compound : (N 0 2 ) 2 0 2 Hg 4- 
2HgO + H 2 0 , which separates as yellow crystals insoluble in water. 
When any of these compounds are gently heated, brown nitrous 
fumes are evolved and red mercuric oxide remains behind. 

Mercuric Chromate is thrown down as a brick-red precipitate 
when potassium chromate is added to a solution of mercuric 
nitrate. 

Mercuric Chloride : HgCl 2 .—This important salt, which is 
commonly called corrosive sublimate , crystallizes from its aqueous 
solution in colourless needles of specific gravity 5-4. It melts to a 
clear liquid at 288°, boils at about 300°, and can be sublimed un¬ 
changed. The salt requires 14 parts of cold water for solution, 
but less than 2 parts of boiling water. It is more easily soluble in 
alcohol and in ether. 

Mercuric chloride may be obtained by dissolving the metal in 
aqua regia, when it crystallizes from the concentrated acid solution. 
It is, however, more commonly prepared in the dry way by sub¬ 
limation. For this purpose a mixture of common salt and mercuric 
sulphate is heated in a glass retort or flask, on a sand-bath, when 
the mercuric chloride thus produced condenses in the upper and 
cooler parts of the vessel as a crystalline cake, which is removed 
by breaking open the vessel when cold. The operation must be 
carried on under a good draught on account of the poisonous nature 
of the vapours of the chloride, some of which always escape into 
the air. The following reaction represents the process :— 

S 0 2 - 0 2 Hg + 2 NaCl = HgCl 2 + S 0 2 ( 0 Na) 2 . 



537 


Mercuric Compounds, 


Mercuric chloride readily parts with one-half of its chlorine 
when heated with substances which easily combine with this ele¬ 
ment, and is converted into mercurous chloride. Thus, an intimate 
mixture of arsenic and mercuric chloride gives arsenious chloride 
when heated (p. 247). 

Corrosive sublimate is an important medicine, but extremely 


poisonous, and can therefore only be given in small doses ( to ^ 
grain, or 4 to 8 milligrammes). Its solution coagulates albumen, 
and possesses powerful antiseptic properties. Wood which has 
been impregnated with a solution of corrosive sublimate is pro¬ 
tected both from decay and from the attacks of insects. It has 
been proposed to saturate the wood of wooden bedsteads with its 
solution, but the process is not to be recommended on account of 
the poisonous nature of the dust given off by the dry wood. 

Caustic soda when added to a solution of mercuric chloride 
gives a yellow precipitate of mercuric oxide, but ammonia when 
added to the same solution behaves differently, and produces a 

white precipitate of the composition : n j Hg 2 > which is to be 


considered as a diammonium dichloride, in which one-half of the 
hydrogen (4 atoms) has been displaced by mercury (2 atoms), 
and is, therefore, dimercuri-diammonium chloride. This substance 
is used in medicine under the name of infusible white precipitate. 

If a mixture of mercuric and ammonium chlorides is precipi¬ 
tated with sodium carbonate, a similar white compound is obtained 
of varying composition, and which melts, when heated, to a yellow¬ 
ish liquid. This substance is called fusible. white precipitate 

When a solution of mercuric chloride is gradually added to a 
boiling mixture of ammonium chloride and ammonia, so that the 
precipitate which is produced redissolves, the solution deposits on 

cooling colourless rhombohedra of the composition : NHjC i[ H 0 
-ie mercuri, diammonium chloride. This compound is closely 
related to the infusible white precipitate, from which it may be 
obtained by boiling with ammonium chloride 


nh 2 ci\ h 

NH.Cl) Hg - 


2 NH 4 C 1 


.NHjCll Hg 
2 NH s C 1 ( s 


Mercuric Bromide : HgBr 2 , resembles the chloride, but is more 
difficultly soluble in water. 


538 Text-Book of Inorganic Chemistry. 

Mercuric Iodide : Hgl 2 .—When solutions of mercuric chloride 
and potassium iodide are mixed, a yellow precipitate of mercuric 
iodide is first produced, which soon turns red, and is soluble either in 
excess of mercuric chloride or of potassium iodide. On cooling its 
warm saturated solution in potassium iodide or in alcohol, the salt 
crystallizes in fine red tetragonal pyramids. It is insoluble in 
water. When heated the salt becomes yellow, then melts and sub¬ 
limes in bright yellow crystals belonging to the rhombic system. 
This yellow modification of mercuric iodide readily passes into the 
red form ; mere contact with the point of a needle is sufficient to 
convert an entire quantity of the yellow crystals into the other 
modification. 

A solution of mercuric iodide in potassium iodide, and made 
alkaline by caustic potash, is used for the detection of minute 
quantities of ammonia or ammonium salts, and is called Nessler's 
solution (p. 387). 


Mercuric Cyanide : HgCy 2 .—Red mercuric oxide readily dis¬ 
solves when digested with dilute hydrocyanic acid, and mercuric 
cyanide crystallizes from the clear solution in brilliant colourless 
quadratic prisms. The crystals contain much enclosed water, but 
none combined as water of crystallization ; the substance is ex¬ 
tremely poisonous. Mercuric cyanide is decomposed on heating 
into mercury and cyanogen, a portion of the latter gas taking the 
form of brown paracyanogen (p. 311). If on digestingmercuric oxide 
with hydrocyanic acid the oxide is in excess, an alkaline solution 
is obtained, from which crystals of a basic cyanide : HgCy 2 + HgO, 
may be separated. 


Mercuric Sulphide, Cinnabar : HgS.—This compound, distin¬ 
guished by its beautiful red colour, is the commonest form in which 
mercury occurs in nature, and is the substance exclusively used for 
the extraction of the metal. 

The mineral cinnabar is sometimes found in small, transparent, 
red crystals, belonging to the hexagonal system, but generally 
occurs in compact dark-red masses. It volatilizes unchanged when 
heated, and in this respect differs from the corresponding red oxide. 
Cinnabar may be obtained artificially by rubbing together a mixture 
of 25 parts of mercury with 4 parts of flowers of sulphur. This 
first produces an amorphous black sulphide, which when heated 


Mercurous Compounds . 539 

in a suitable vessel yields a sublimate of cinnabar as a translucent, 
red crystalline mass of fibrous structure. 

Mercuric sulphide is produced in the wet way by leading sulphu¬ 
retted hydrogen into a solution of a mercuric salt. A black preci¬ 
pitate is then produced, which has the same composition as cinnabar, 
and into which it may be converted by sublimation. When 
the sulphuretted hydrogen is led into the mercuric salt a white 
precipitate is first thrown dowh, which soon becomes yellow, then 
red, and finally black. These intermediate substances are double 
compounds, in varying proportions, of mercuric sulphide and the 
mercuric compound used for the precipitation. 

Mercuric sulphide is insoluble both in nitric acid and in 
hydrochloric acid, and only dissolves in aqua regia. This pro¬ 
perty enables us to separate it from the sulphides of other metals 

_ e.g, copper—which dissolve in nitric acid. The sulphide in the 

form of powdered cinnabar is the valuable pigment vermilion , 
which is prepared either by grinding sublimed cinnabar or by 
various wet processes. The colour is improved by continued rub¬ 
bing or by allowing the powder to stand for some time in contact 
with a solution of potassium pentasulphide at a temperature of 50°. 


Mercurous Compounds. 

Mercurous Oxide : Hg.O.—This compound is much less stable 
than mercuric oxide; it may be obtained as a black amorphous 
precipitate by mixing a solution of mercurous nitrate with a slight 
excess of caustic'soda. When rubbed, or when heated to ioo°, it 
decomposes into mercuric oxide and free mercury. The same 
change is also produced by the action of light, and it must, there¬ 
fore be filtered, washed, and dried in the dark. Undecomposed 
mercurous oxide dissolves completely in dilute nitric acid. 


Mercurous Nitrate: (N 0 ,), 0 ,Hg, + 2 H, 0 ^-If mercury is 
covered with a layer of moderately dilute nitric acid and allowed 
to stand for some time at the ordinary temperature, colourless 
transparent crystals gradually separate out in the form of rhombic 
plates This salt is mercurous nitrate. The crystals easily dis- 
solve'in the mother-liquor when warmed, and reseparate when 
the solution again cools. The salt is decomposed by water, with 
!he separation of a yellow basic salt: (N 0 2 ) 2 0 2 Hg 2 + Hg a O + H 2 0 . 


54 ° Text-Book of Inorganic Chemistry . 

A second basic salt, containing less mercurous oxide, is obtained 
by warming nitric acid with excess of mercury ; it crystallizes 
from its hot saturated solution in transparent colourless prisms. 

On the addition of ammonia to a solution of mercurous nitrate 
a black precipitate, resembling mercurous oxide, is obtained. This 
substance contains, however, the elements of ammonia and has a 
varying composition ; its chief constituent is probably a compound 

of the composition : N0 2 -0NH 2 j Hg 4 — dimercuroso-diammonium 

nitrate — i.e. diammonium nitrate in which four atoms of hydrogen 
are displaced by four monad atoms of mercury. It was formerly 
used in medicine under the name of Mercurius solubilis Hahne- 
manni. 

Mercurous Chromate : Cr 0 2 - 0 2 Hg 2 , is thrown down as a red 
precipitate on adding a solution of potassium chromate or dichro¬ 
mate to one of mercurous nitrate ; it is quite insoluble in water, and 
is used for the separation and quantitative estimation of chromic 
acid. The dry salt decomposes when heated into mercury and 
oxygen, which are given off, and green chromic oxide, which remains 
behind. From the weight of chromic oxide so obtained that of the 
chromic acid precipitated may be easily calculated. 

Mercurous Chloride, Calomel : Hg 2 Cl 2 .—This salt, which like 
corrosive sublimate is a valuable medicine, is distinguished by its 
complete insolubility in water. It is found in nature as the mineral 
horn-quicksilver , and may be prepared artificially either in the wet 
way (by precipitation) or in the dry way (by sublimation). By the 
former method, it is obtained as a white curdy precipitate, when 
hydrochloric acid or a soluble chloride is added to a solution of 
mercurous nitrate. This precipitate when washed and dried forms 
a heavy white powder. Calomel is also obtained in tlie dry way 
when a mixture of mercuric chloride and mercury is heated. In 
grinding the two substances together it is well to add a little 
alcohol to prevent the formation of any dust of the poisonous 
mercuric chloride. Thin glass flasks are then about one-third 
filled with this mixture, and heated on a sand-bath ; vapours of 
mercury and mercuric chloride are thus formed, and unite together 
to produce mercurous chloride, which sublimes and condenses in 
the upper and cooler parts of the flasks. The operation must be 
conducted under a hood with a good draught on account of the 


Amalgams. 54 1 

extremely poisonous nature of the vapours of mercuric chloride, 
some of which always escape from the flasks. 

This process is not, however, used for the manufacture of calo¬ 
mel, but a mixture of mercuric sulphate and common salt (which 
yields mercuric chloride when heated) is ground up with an equal 
quantity of mercury to that already contained in the sulphate, and 
then heated in glass or earthenware vessels. Mercurous chloride 
then sublimes and sodium sulphate remains behind. 

Calomel prepared in either of these ways always contains more 
or less corrosive sublimate, which may be extracted from it by 
digesting the powdered substance with warm alcohol. This puri¬ 
fication is especially necessary if the calomel is to be employed for 
medicinal purposes. Calomel is much less poisonous than corro¬ 
sive sublimate, and can therefore be given in larger doses. 

Mercurous chloride when treated with ammonia becomes con¬ 
verted into a black powder of the composition : ( Hg 4 — di- 

mercuroso-diammoniam chloride , and in this way may be distin¬ 
guished from silver chloride, which dissolves in ammonia. 

Mercurous Bromide : Hg 2 Br 2 , closely resembles the chloride in 
its properties. 

Mercurous Iodide : Hg 2 I 2 , is thrown down as a dirty green pre¬ 
cipitate when a solution of potassium iodide is mixed with one of 
mercurous nitrate : it may also be obtained by rubbing together 
mercury and iodine in the correct proportions. The compound 
turns black in the light, and when heated is decomposed into mer¬ 
curic iodide and mercury. 

Mercurous Cyanide has not yet been prepared. 

Mercurous Sulphide : Hg 2 S, is possibly produced when sulphu¬ 
retted hydrogen is led into a solution of mercurous nitrate ; it is, 
however, very unstable, and soon decomposes into mercuric sul¬ 
phide and mercury. 

Amalgams. 

Mercury readily unites with most metals, and produces alloys 
—generally known as amalgams. Often mere contact at the ordi¬ 
nary temperature suffices to produce an amalgam, as in the case 
of silver, gold, and the alkali metals. Most amalgams may be 


54 2 Text-Book of Inorganic Chemistry. 

obtained in the solid crystalline form, and are all soluble in ex¬ 
cess of mercury. An amalgam of iron can only be obtained with 
difficulty, and iron vessels are therefore employed to contain the 
metal. 

Sodium amalgam may be easily obtained as a solid crystalline 
mass by adding 4 parts of sodium, in small quantities at a time, to 
100 parts of warm mercury contained in a mortar. As each small 
piece of sodium is pressed under the mercury with the pestle, the 
two metals unite with a hissing noise, and a sufficient evolution of 
heat to volatilize some of the mercury. The amalgam remains 
liquid as long as it is warm, but becomes a crystalline solid when 
cold. Mercury unites still more energetically with potassium. 

A piece of gold when dipped into mercury becomes at once 
covered with a white layer of gold amalgam, and soon completely 
dissolves. Silver also easily dissolves in an excess of mercury. 
The readiness with which these two metals unite with mercury is 
utilized to separate them from their ores. These ores when shaken 
with mercury yield amalgams, from which the mercury can be after¬ 
wards distilled, leaving silver or gold behind (amalgamation pro¬ 
cess). One of the most useful amalgams is that of tin, which is 
used for silvering ordinary looking-glasses (p. 522). 

Metallic mercury is used for a number of technical and scientific 
purposes—for example, in the construction of barometers and ther¬ 
mometers, and for the mercury air-pump (Sprengel’s pump). It is 
also used in the laboratory to collect gases which are dissolved by 
water but which do not act upon mercury— eg. ammonia, hydro¬ 
chloric acid. Chlorine cannot be collected over mercury, as the two 
substances at once unite with one another and produce a chloride. 


Detection of IVIercury Compounds. 

All solid compounds of mercury when mixed with quick-lime or 
dry sodium carbonate and heated in a small tube are decomposed, 
and yield metallic mercury, which condenses in the cooler parts of 
the tube in minute glistening globules. 

Soluble mercuric salts may be recognized by the yellow pre¬ 
cipitate of mercuric oxide produced by caustic soda and the white 
precipitate formed on the addition of ammonia. Black mercuric 
sulphide, which is produced when a stream of sulphuretted hydrogen 
is passed through a solution of a mercuric salt, is distinguished 
from other black sulphides produced in the same way (eg. those 


Silver. 


543 


of copper, lead, bismuth) by remaining unchanged when warmed 
with nitric acid. The soluble mercurous salts, of which only one or 
two are known, give a white precipitate of mercurous chloride with 
hydrochloric acid, which blackens on the addition of ammonia. 
Caustic soda and ammonia, as well as ammonium or sodium 
carbonate, all produce black precipitates with a solution of mer¬ 
curous nitrate. If a solution of stannous chloride is added to a 
solution of a mercuric or mercurous salt, a white precipitate of 
mercurous chloride is produced in both cases, in the former with 
the simultaneous production of stannic chloride. And if this pre¬ 
cipitate of mercurous chloride is mixed with a further quantity of 
stannous chloride and some hydrochloric acid, it is reduced to a 
grey powder of metallic mercury, the particles of which unite 
when boiled and produce small globules of the metal. 


SILVER. 

Chemical Symbol : Ag .—Atomic Weight'. 108. 

Silver is tolerably widely distributed in nature, but is seldom 
found in large quantities. It occurs as native silver crystallized in 
regular octahedra, and united with chlorine as horn-silver (AgCl), 
as well as with the other halogens bromine and iodine, or even 
with all three together : it is also occasionally found mixed with 
mercury as amalgam. More important than these are the com¬ 
pounds of silver with sulphur, silver-glance (Ag 2 S), and those con¬ 
taining silver sulphide united with other sulphides (eg. those of 
arsenic, antimony, and copper). These double sulphides form the 
minerals dark-red silver ore or pyrargyrite (3Ag 2 S,Sb 2 S 3 ), light- 
red silver ore or proustite (3Ag 2 S,As 2 S 3 ), and silver copper-glance 
(Ag 0 S,Cu 2 S). 

A considerable quantity of silver is extracted from argentiferous 
lead by the process of cupellation, which has been already de¬ 
scribed under lead (p. 500)- From ores containing little lead the 
silver is extracted by the process of amalgamation ,, which varies in 
different countries according to the nature of the ore and according 
as fuel is cheap or dear. The rich ores of Nevada and Colorado 
in the United States are roasted with common salt, by which the 
silver loses its sulphur and is converted into silver chloride. The 
roasted ores are then mixed with water and scrap iron and rotated 



544 Text-Book of Inorganic Chemistry. 

in barrels for some time, when the iron gradually reduces the silver 
to the metallic state. As soon as this change has been produced, 
mercury is introduced into the barrels, and the silver amalgam so 
obtained run off, dried, and distilled. Mercury then passes over 
into a condenser and metallic silver remains behind. In Mexico 
and Chili, where fuel is scarce and dear, the silver is converted into 
chloride by a mechanical process which consists in mixing the 
ground ore with common salt and salts of copper and iron by the 
treading of mules ; mercury in excess is next added, part of which 
reduces the silver chloride with formation of calomel, the remainder 
forming an amalgam which is afterwards distilled. Various wet 
processes are also used for the extraction of silver from very poor 
ores. 

Pure silver may be obtained by precipitating a solution of a 
silver salt with hydrochloric acid, and fusing the pure dry silver 
chloride so produced with dry sodium carbonate in an unglazed 
porcelain crucible. When cold, the silver is found at the bottom of 
the crucible covered with a layer of fused sodium chloride. 

Silver is a white, lustrous metal, with a specific gravity of io'6. 
It is soft, and so malleable that it can be beaten out into extremely 
thin leaves (silver leaf). Silver melts at a lower temperature than 
copper, and although it does not oxidize in the air, possesses the 
property of dissolving oxygen when in the liquid state. This oxy¬ 
gen, which is only dissolved in the silver not chemically combined 
with it, is again given off when the silver solidifies, producing the 
phenomenon known as the spitting of silver. Pure silver is said 
to absorb as much as 22 times its volume of oxygen. The property 
is only possessed by the pure metal: the addition of a small 
quantity of lead prevents the absorption of the gas and the subse¬ 
quent spitting. 

Silver combines far more readily with sulphur than with oxygen. 
Sulphuretted hydrogen is decomposed by silver, and the metal be¬ 
comes covered with a thin brown or black layer of silver sulphide. 
Hydrochloric acid does not attack the metal, but it is readily dis¬ 
solved by nitric acid with evolution of brown nitrous fumes, and 
formation of silver nitrate. Concentrated sulphuric acid dissolves 
silver in the same way as copper and mercury, producing silver 
sulphate and sulphurous anhydride. Silver is precipitated from 
solutions of its salts by other more positive metals— eg. zinc, iron, 
and copper—and also by certain reducing substances, such as 
ammonium sulphite, grape sugar, aldehyde, and tartaric acid. 


Silver. 


545 


Some of these substances cause a deposit of the silver as a uniform 
coating on the walls of the vessel used for the experiment, and pro¬ 
duce a mirror of metallic silver. 

All commercial silver, including silver coins, always contains 
from 5 to io per cent of copper. To prepare pure silver from this 
alloy, the metal may be dissolved in nitric acid, the solution filtered 
from a trace of gold which remains behind as a black powder, and 
the filtrate precipitated with hydrochloric acid or a solution of com¬ 
mon salt. The silver chloride is then washed, dried, mixed with 
dry sodium carbonate, and heated in an unglazed crucible. Sodium 
chloride and silver carbonate are then produced, and the latter com¬ 
pound is decomposed at a red-heat into carbonic acid, oxygen, and 
silver. After cooling, a button of pure silver is found at the bottom 
of the crucible covered with a layer of fused sodium chloride. 


Compounds of Silver. 

Silver is one of the few metals which exists as a monad element 
in nearly all its compounds. In this respect it resembles the 
alkali metals and thallium, and like sodium it also forms a per¬ 
oxide, containing the metal as a dyad. Its sulphate unites with 
aluminium sulphate when heated in sealed tubes, and produces 
silver alum, isortiorphous with the alums of potassium and 
thallium. 

Silver Oxide : Ag 2 0 , is precipitated as a brown amorphous 
powder when silver nitrate is mixed with pure caustic soda or some 
other soluble base. As ordinary caustic soda always contains 
sodium chloride, the silver oxide obtained by it is mixed with silver 
chloride, and it is therefore preferable to use baryta-water. Silver 
oxide is not quite insoluble in water ; it imparts a metallic taste 
and a faint alkaline reaction to it. When exposed to the light it 
slowly decomposes at the ordinary temperature into silver and 
oxygen—a change which takes place rapidly when the compound 
is heated. Hydrogen reduces it completely even at ioo°. Silver 
oxide unites readily with all acids, but few of its salts are soluble. 

If freshly precipitated, silver oxide is digested with strong am¬ 
monia ; it is converted into a black powder, which may also be 
obtained by adding caustic potash to an ammoniacal solution of 
silver nitrate. This substance is one of the most unstable com¬ 
pounds known : it explodes violently when slightly pressed, even 

N N 


546 Text-Book of Inorganic Chemistry. 

when moist. It is known as Berthollef s fulminating silver, but we 
are quite ignorant of its composition. Possibly it is silver amide : 
NH 2 Ag, or silver nitride : NAg 3 . 

Silver Peroxide : AgO or Ag 2 0 ,.—This little-studied compound 
is obtained in the form of black shining octahedra when ozone acts 
on a moist silver plate, or when an electric current is passed through 
a strong solution of silver nitrate. It is then deposited on the 
positive pole, for which it is best to employ a stout platinum wire. 
The compound is less stable than silver oxide ; it decomposes when 
heated over ioo° into silver oxide and oxygen, and when inflam¬ 
mable substances, such as sulphur or phosphorus, are brought into 
contact with it, they are oxidized with an explosion. When warmed 
with nitric acid it produces silver nitrate, and oxygen is evolved. 
Cold nitric acid dissolves it and produces a brown liquid. 

Silver Sulphate :.S 0 2 ( 0 Ag) 3 .—When silver is warmed with con¬ 
centrated sulphuric acid, the metal gradually dissolves, with evolution 
of sulphurous acid, and silver sulphate is deposited as small glisten¬ 
ing crystals. The salt is difficultly soluble in water, and may there¬ 
fore be prepared by adding a solution of sodium sulphate to one of 
silver nitrate. As the salt is far more soluble in hot than in cold 
water, it is deposited when the above solutions are made hot, mixed, 
and the mixture allowed to cool. It crystallizes in small rhombic 
prisms, which are isomorphous with anhydrous sodium sulphate. 
At a low red-heat, silver sulphate melts unchanged, but at higher 
temperatures it is decomposed into silver, sulphurous and sulphuric 
anhydrides, and oxygen. 

Silver Nitrate : N 0 2 - OAg.—This, the most important salt of 
silver, is obtained by dissolving pure silver in nitric acid and evapo¬ 
rating down the acid liquid. It crystallizes in large colourless 
rhombic plates, is easily soluble in water, and the solution reacts 
neutral. Alcohol also dissolves considerable quantities of the salt. 
Silver nitrate melts at about 220°, and solidifies on cooling to a 
crystalline mass. At a higher temperature it is converted into 
silver nitrite, which when more strongly heated is completely de¬ 
composed, metallic silver remaining behind. 

If silver nitrate is carefully melted at the lowest possible 
temperature and poured into moulds, similar to those used for 
caustic potash (p. 348), it is obtained in the form of white solid 


Silver Nitrate. 


54 7 


sticks, which are used as a caustic in surgery under the name of 
lunar caustic. Silver nitrate destroys those bodies which cause the 
putrefaction and decomposition of organic substances, and also 
coagulates albumen ; hence the spots touched with the caustic 
become covered with a hard crust of coagulated albumen and fibrin, 
which soon blackens, especially when exposed to the light. The 
blackening is due to finely divided metallic silver reduced from the 
nitrate by the organic matter. A solution of silver nitrate also 
blackens when exposed to the light in contact with organic sub¬ 
stances, and such a solution mixed with gum water is used for 
marking linen and cotton articles. Characters written with this 
marking-ink gradually blacken when warmed or when exposed to 
light, and the reduced silver adheres so firmly to the fibres of the 
stuff, that it is not removed by washing with soap and soda. Such 
marks may, however, be obliterated by moistening them with a solu¬ 
tion of potassium cyanide and then very carefully washing with water, 
on account of the extremely poisonous character of this compound. 
Potassium cyanide dissolves finely divided silver in the presence 
of water, producing soluble potassium-silver cyanide with some 
caustic potash, and free hydrogen :— 

2KCy + Ag + H 2 0 = KCy,AgCy + KOH + H. 

Silver nitrate has many applications in the arts. Large quanti¬ 
ties are used in photography for the production of the halogen 
compounds of silver which are readily decomposed by light. 

Silver Phosphate : PO(OAg) 3 , is produced as a bright yellow 
precipitate by mixing a solution of common (monacid) sodium 
phosphate with one of silver nitrate :— 

PO(ONa) 2 (OH) + 3N0 2 -OAg = PO(OAg) s + 2 N 0 2 0 Na 

+ no 2 -oh. 

The filtered solution then reacts acid from free nitric acid, and 
contains some,silver dissolved as acid silver phosphate. 

Silver Arsenite : As(OAg) 3 , closely resembles the yellow phos¬ 
phate. Silver Arsenate : AsO(OAg) 3 , is a reddish-brown powder, 
and, like the phosphate and arsenite, easily dissolves in nitric acid. 

Silver Carbonate : CO(OAg) 2 , is thrown down on mixing a 
solution of silver nitrate with one of sodium carbonate as a 
yellowish-white precipitate, which quickly darkens in the light. 

N N 2 


548 


Text-Book of Inorganic Chemistry. 


The dry salt loses its carbonic acid at a temperature slightly above 

IOO°. 

Silver Chromate-. Cr 0 2 ( 0 Ag) 2> is a dark-red precipitate ob¬ 
tained by mixing solutions of potassium chromate and silver 
nitrate. 


Silver Chloride, Bromide, and Iodide. 

These three halogen compounds of silver closely resemble one 
another, especially in their chemical propert.es. They all occur m 
nature, and often associated with one another. They are obtained 
as curdy precipitates of a white or pale yellow colour, and insoluble 
in dilute acids when the corresponding halogen acid or haloid salt 
is added to a solution of silver nitrate. Silver chloride is of a pure 
white colour, the bromide is faintly yellow, and the iodide has a 
more decided yellow tint. They all melt when heated without e- 
composition, and are reduced to the metallic state by zinc when 
this metal is dipped into one of the fused salts or when zinc is 
added to the salt suspended in water containing hydrochloric acid. 
All three dissolve easily in an aqueous solution of sodium thiosul¬ 
phate or of potassium cyanide, but are distinguished from one 
another by their solubility in ammonia. Silver Chloride . AgU, 
when freshly precipitated, easily dissolves in a small ^tantity of 
ammonia; Silver Bromide -. AgBr dissolves only with difficulty and 

requires a large quantity of the reagent; while Silver Iodide : Agl, 
is almost insoluble in ammonia, but is dissolved by P°‘ ass > u ™ 
iodide, forming a soluble double salt of the composition: KI,Agl, 
which crystallizes out when the solution is evaporated down. 
Silver is scarcely attacked by hydrochloric or hydrobromic acid, 
but strong hydriodic acid easily converts the metal into silver 

iodide with evolution of hydrogen. 

Silver bromide and iodide are less stable compounds than silver 
chloride Both the former are decomposed when heated in a stream 
of chlorine; silver chloride is then produced, with separation of bro¬ 
mine or iodine respectively. But if a mixed solution of potassium 
iodide and chloride is added to a small quantity of silver nitrate, the 
compound first produced is not silver chloride but silver iodide. 
In the same manner, silver chloride is completely decomposed into 
silver bromide when boiled with a solution of potassium bromide, 
and both silver chloride and bromide become converted into silver 
iodide when treated with potassium iodide. From these results it 


Halogen Compounds of Silver . 549 

has sometimes been imagined that the affinity of silver for chlorine 
in an aqueous solution is less than for bromine ; but if we consider 


the three equations 

which 

represent these changes :— 

AgCl 

+ 

KBr = AgBr + KC 1 

AgBr 

+ 

KI = Agl + KBr 

AgCl 

+ 

KI = Agl + KG 1 


we see that potassium chloride is formed in the first and third 
equations and potassium bromide in the second, and potassium 
chloride is a much more stable compound than either the iodide or 
or the bromide, while the bromide is more stable than the iodide. 
The cause of the reactions is therefore due not to a stronger affinity 
of silver for bromine or iodine than for chlorine, but to a stronger 
affinity of potassium for chlorine than for bromine and iodine, and 

for bromine than for iodine. , 

The sensibility of the halogen compounds of silver to light, 
under the influence of which, in the presence of reducing agents, 
they are reduced to sub-salts and then to metallic silver, makes 
them very suitable for the production of photographs. The com¬ 
pounds, usually silver bromide or iodide, are suspended in some 
suitable medium (collodion or gelatine), and glass plates coated 
with the mixture. These plates are then exposed in the camera to 
the light reflected from the object of which a picture is required, 
when the silver iodide or bromide is more or less reduced according 
to the intensity of the light falling upon it, and according to the 
length of time which the plate is exposed. On removing the plate 
from the camera no change can be seen, and no metallic silver has 
yet been produced : the light has only formed a sub-bromide or 
iodide of the same colour as the original salt. But if the exposed 
plate is brought into a solution of some reducing substance, such 
as pyrogallic acid or ferrous oxalate, the reducing action which the 
light has begun is carried further and the sub-bromide or -iodide 
is now converted into black metallic silver. This process is called 
developing;. In order to preserve this picture and to prevent the 
undecomposed silver salt from blackening when exposed to the 
light, it must now be fixed— a process which consists in dissolving 
out the undecomposed bromide or iodide by a solution of sodium 
thiosulphate or potassium cyanide. The picture so obtained shows 
the bright parts of the object dark and the dark parts light, and is 
called a negative ; an ordinary photograph on paper is obtained 
from this by placing the glass negative over a sheet of paper con- 


550 Text-Book of Inorganic Chemistry. 

taining silver chloride in its pores and then exposing to the light. 
The print so obtained is then treated with various solutions con¬ 
taining gold, to make the colour more pleasing, and finally fixed 
with sodium thiosulphate. Not only does the intensity of the 
light influence its reducing action on the halogen compounds, but 
light of different colour produces very different effects. Thus the 
blue and violet end of the spectrum has a far more powerful re¬ 
ducing action than the green, yellow, and red ; and the invisible 
rays which lie beyond the visible violet (the ultra-violet or actinic 
rays) act most powerfully of all. Hence all photographic processes 
must be conducted in a room which is only illuminated with yellow 
or red light. 

Silver Fluoride : AgF, is distinguished from the other halogen 
compounds of silver by its solubility in water. It may be easily 
obtained by dissolving silver carbonate in hydrofluoric acid, and 
crystallizes from its aqueous solution with two molecules of water. 

Silver Cyanide : AgCy, is a white curdy precipitate produced 
by carefully adding a solution of potassium cyanide to silver nitrate. 
It is insoluble in water and dilute acids, but easily dissolves in 
ammonia. An excess of potassium cyanide also dissolves it, and 
produces a soluble double cyanide of the composition : KCy,AgCy, 
which may be easily obtained in the crystalline form. Unlike the 
chloride, bromide, and iodide, it is decomposed when heated, 
yielding cyanogen gas. The grey substance which then remains 
behind is a mixture of metallic silver and paracyanogen. The 
soluble compound : KCy,AgCy, is largely used in silver plating. 
The silver which is deposited at the negative pole when an electric 
current is passed through a solution of this double cyanide is in a 
more compact form, and adheres firmer to the substance to be 
plated than that deposited in a similar manner from any other 
solution of silver. 

Silver Sulphide : Ag. 2 S.—Silver possesses a strong affinity for 
sulphur, and even decomposes sulphuretted hydrogen in the presence 
of air, but remains unchanged in the pure dry gas. The compound 
may be obtained by melting silver and sulphur together, or by pre- 
ciditating a salt of silver with sulphuretted hydrogen. When 
prepared by the latter method it is a black powder easily soluble 
in nitric acid. Silver sulphide occurs in the mineral kingdom as 


Alloys of Silver. 55 1 

silver-glance, and in combination with the sulphides of arsenic or 
antimony as light and dark red silver ore respectively. 


Alloys of Silver. 

The numerous articles which are made of silver do not consist 
of the pure metal, but of an alloy of copper and silver. The pure 
metal is too soft for ordinary use and would wear away rapidly. 
But the addition of a few parts per cent, of copper considerably 
increases it in hardness and makes it more suitable for ordinary 
purposes. The two metals fuse together in all proportions, and 
those alloys which only contain 20 per cent, of silver have a 
reddish colour, while those containing 80 to 90 per cent, are white. 
English silver coins contain 92-5 per cent, of silver and r 5 per 
cent, of copper, while the silver coinage of France, Germany, 
and Austria, consists of an alloy of 90 parts of silver with 10 parts 
of copper. The percentage of silver contained m a given alloy is 
usually expressed in parts per thousand or per mills. Thus the 
fineness of the silver used in England for coinage would be 92 5, and 
that used in France and Germany for the same purpose would be 
900. The fineness of the silver alloys which are to be employed for 
different purposes is regulated by law in some countries. 

The quantity of silver contained in an alloy with copper was 
formerly exclusively determined by cupellation. In this process 
the sample of silver is mixed with lead and heated on a small flat 
crucible made of bone-ash (the cupel) in a current of air ; the lead 
and the copper are then oxidized, and their oxides are absorbe 
by the porous bone-ash of the cupel. Finally, there remains behind 
a button of pure silver, from the weight of which the fineness of 
the sample may be easily calculated. A quicker and more accu ^ at ® 
process is the wet way introduced by Gay-Lussac. In this method 
the silver is dissolved in nitric acid and a normal solution of 
common salt added until the whole of the silver is precipitated as 
silver chloride. For the details‘of this method the student must 
consult a text-book of analytical chemistry. 


Detection of Silver Compounds. 

The reduction of any compound of silver when mixed with 
sodium carbonate and heated on charcoal before the blowpipe, and 
the production of a white globule of metallic silver which does not 
oxidize in the air, the precipitation of black silver sulphide when 


552 Text-Book of Inorganic Chemistry. 

sulphuretted hydrogen is led through an acid solution of a silver 
salt, and, above all, the formation of silver chloride by hydro¬ 
chloric acid, make it very easy to detect the presence of silver and 
to separate it from other metals. Silver chloride may be distin¬ 
guished from lead chloride by the solubility of the latter salt in hot 
water, and from mercurous chloride by the action of ammonia. 
This dissolves the silver chloride, but converts the mercurous salt 
into a black compound (p. 543). 


GOLD. 

Chemical Symbol : Au .—Atomic Weight : 197. 

Gold is chiefly found in nature in the free state, sometimes pure 
and well crystallized, but usually containing silver. It occurs in 
quartz rocks or in alluvial quartz sand which has been deposited by 
the disintegration of the older rocks. These auriferous sands, 
which often cover a considerable extent of country, are washed by 
suitable arrangements, on a small or large scale, when the lighter 
particles of sand and mud are carried away and the heavier particles 
of gold remain behind. The gold is then brought into the market 
in this form, or if in a finely divided state is extracted with mercury, 
which is afterwards removed by distillation. Gold is now chiefly 
obtained from California and Australia, but is found in smaller 
quantities in other parts— e.g. in the Ural Mountains, in Hungary 
and Transylvania, and in Africa. Small quantities of gold have also 
been found in Cornwall and Wales. Chemical compounds of gold 
rarely occur in nature —sylvanite or graphic tellurium , found in 
Transylvania, is a compound of gold and tellurium. Many silver 
ores and some argentiferous lead ores, as well as some specimens 
of copper-pyrites, contain small quantities of gold. In the silver 
ores the gold is extracted with the silver, and the two metals are 
afterwards separated by nitric acid. 

Gold belongs to the few metals which have a decided tint, and 
is distinguished from all others by its yellow colour. It crystallizes 
in the regular system, is very soft, considerably more so than silver, 
and very malleable. In extremely thin leaves (gold-leaf), it trans¬ 
mits bluish-green light. Gold melts more difficultly than silver, 
but somewhat more easily than copper, and volatilizes at very high 
temperatures. Its specific gravity is 19*3. 



Gold. 


553 


Gold does not directly unite with oxygen at any temperature, 
but easily combines with chlorine. Hydrochloric, sulphuric, or 
nitric acid has no action on the metal, but aqua regia dissolves gold 
and produces a yellow solution of auric chloride. From this solu¬ 
tion the metal may be again precipitated by various reducing agents, 
such as ferrous sulphate, oxalic acid, sulphurous acid, or metallic 
iron. The best substance to employ for the precipitation of gold 
from a solution of its chloride is a cold saturated solution of ferrous 
sulphate, which is then oxidized to a mixture of ferric sulphate and 
ferric chloride. The gold is thrown down as a brown powder, 
which acquires the colour and lustre of the metal when rubbed. 
This method is used to separate pure gold from a solution con¬ 
taining copper and other metals. 

Compounds of Gold. 

Gold forms two series of compounds in which the element exists, 
as a triad and as a monad respectively, and which are distinguished 
as the auric and aurous compounds. The metal only possesses 
weak affinities for other elements, and its two oxides are such weak 
bases that their salts are hardly known. The most important auric 
compound is 

Auric Chloride : AuC 1 3 , which is obtained by dissolving pure 
gold in aqua regia and remains as a reddish brown deliquescent 
mass when the solution is evaporated to dryness. It not only dis¬ 
solves in water, but also in alcohol and ether, and to such an extent 
in the latter liquid that the chloride is abstracted from its aqueous 
solution when this is shaken with ether. A solution of auric chlo¬ 
ride produces a purple stain of reduced gold when a drop is allowed 
to fall on the skin, and gives a purple precipitate when mixed with 
stannous chloride—see Purple of Cassius (p. 55 ^)* When heated, 
the salt decomposes into chlorine and aurous chloiide. 

Auric chloride forms soluble crystalline double salts with hy¬ 
drochloric acid and soluble chlorides. The hydrochloric acid 
compound: AuC 1 3 ,HC 1 + 3 H 2 0 , or HAuC 1 4 + 3 H 2 0 , is deposited 
from a hot acid solution of auric chloride in long yellow needles. 
On mixing a solution of auric chloride with one of the chlorides of 
potassium, sodium, or ammonium, and evaporating down, crystals 
of the compounds : 2 (AuC1 3 ,KC1) + H 2 0 , or 2 KAuC1 4 + H 2 0 ; AuCL, 
NaCl + 2H„0, or NaAuCl 4 + 2H q O, and 2 AuC1 3 ,NH 4 C1 + 5 H 2 0 , or 
2 NH 4 AuC1 4 + 5H,0 are obtained. The sodium compound is much 


554 Text-Book of Inorganic Chemistry . 

less deliquescent than pure auric chloride, and is the form in which 
this salt is usually brought into trade. It is sold as ?ton-deliquescent 
chloride of gold in 15 grain or 1 gramme tubes. 

Auric Oxide : Au 2 0 3 , is obtained by digesting a solution of 
auric chloride with an excess of magnesia. The brown powder so 
produced consists of a compound of auric oxide and magnesia 
(magnesium aurate) with the excess of magnesia. When this 
powder is treated with strong nitric acid, the magnesia is all dis¬ 
solved and auric oxide remains behind as a brown powder. It is 
decomposed by light or heat into gold and oxygen. If the brown 
compound of auric oxide and magnesia is digested with dilute in¬ 
stead of concentrated nitric acid, auric hydrate : Au(OH) 3 , remains 
behind as a reddish yellow powder. 

Auric oxide is so weak a base that neither it nor the hydrate 
can unite with oxyacids to form salts. On the other hand, it can 
combine with bases and form salts in which it plays the part of an 
acid. Hence auric oxide may be considered as the anhydride of 
an acid— auric acid. 

Potassium Aurate : AuO • OK + 3 H 2 0 .—Auric oxide or hydrate 
easily dissolves in caustic potash, and when the solution is eva¬ 
porated, finally in a vacuum, potassium aurate separates in bright 
yellow needles. Its solution reacts alkaline and gives precipitates 
of insoluble aurates when mixed with the salts of other metals. 

If a solution of auric chloride is digested with an excess of 
ammonia a yellowish brown powder is produced, which violently 
explodes when slightly warmed or rubbed. The composition of 
this fulminating gold is unknown. 

Auric Cyanide: AuCy 3 + 3 H 2 0 .—Nearly all the cyanides of 
the heavy metals are insoluble in water, but mercuric cyanide and 
auric cyanide are exceptions to this rule. If a solution of auric 
chloride is mixed with a hot concentrated solution of potassium 
cyanide, the compound potassium-auric cyanide : KCy,AuCy 3 + 
H 2 0 , or KAuCy 4 + H 2 0 , crystallizes out on cooling in colourless 
plates. This substance, when dissolved in water, gives a white 
curdy precipitate with a solution of silver nitrate of the corre¬ 
sponding silver salt. If this silver-auric cyanide is carefully 
washed, mixed with water, and then a quantity of hydrochloric 
acid added insufficient to produce complete decomposition, silver 
chloride, hydrocyanic acid, and auric cyanide are produced, and 


555 


Compounds of Gold. 

when the clear liquid is evaporated in a vacuum over sulphuric 
acid this salt remains behind as a white crystalline mass. By 
recrystallizing from its alcoholic solution, the auric cyanide may be 
obtained as large crystalline tablets or plates. The salt is easily 
decomposed when heated. 

Auric Sulphide : Au 2 S 3 (?).-—It is still doubtful whether the 
black precipitate which is produced when sulphuretted hydrogen 
is led through a cold solution of auric chloride consists of auric 
sulphide, or is a compound of aurous and auric sulphides. The 
precipitate easily dissolves in solutions of alkaline sulphides. If 
gold is fused with potassium polysulphide (liver of sulphur), a 
double sulphide of potassium and gold is formed, which dissolves 
when the cold mass is digested with water. From this solution 
acids precipitate a sulphide of gold. 

The following more important aurous compounds may be de¬ 
scribed. 

Aurous Chloride : AuCl.—Auric chloride when heated decom¬ 
poses into chlorine and aurous chloride or gold, according to the 
temperature employed. If the temperature does not exceed 2oo° ? 
aurous chloride remains behind as a dirty-white powder which is 
insoluble in water. It is decomposed into auric chloride and 
metallic gold when boiled with water. 

Aurous Oxide : Au 2 0.-Aurous chloride when digested with 
caustic potash is converted into a dark-violet coloured powder of 
aurous oxide. The compound is insoluble in water, and is not 
changed by sulphuric or nitric acid. Hydrochloric acid decom¬ 
poses it into auric chloride and gold. Salts of aurous oxide are 
not known if we except the crystalline salt which separates from a 
mixed solution of auric chloride and sodium thiosulphate on the 
addition of alcohol, and which may be a double salt of aurous and 
sodium thiosulphates. 

Aurous cyanide : AuCy.-Fulminating gold readily dissolves 
in a warm solution of potassium cyanide, and the solution deposits 
large prisms of potassium-aurous cyanide : KCy,AuCy or KAuCy 2 , 
on cooling. Hydrochloric acid separates aurous cyanide from this 
salt, and the cyanide may be obtained by evaporating to dryness 
and extracting with water, when it remains as a yellow crystalline 
powder insoluble in water. Potassium cyanide easily re-dissolves 
it it is decomposed when heated into gold and cyanogen. 


556 Text-Book of Inorganic Chemistry. 

Purple of Cassius is a compound of gold, tin, and oxygen, of 
varying and, therefore, uncertain composition. It is thrown down 
as a beautiful purple precipitate by adding a mixture of stannous 
and stannic chlorides to a dilute solution of auric chloride. If 
concentrated solutions of these salts are mixed, the precipitate 
possesses a brown colour. The compound may also be prepared 
by acting on an alloy of gold, silver, and tin with nitric acid. The 
silver then dissolves and purple of Cassius remains behind. Purple 
of Cassius is used in painting on porcelain and glass to produce a 
beautiful red tint. 


Alloys of Gold. 

Pure gold is used for the manufacture of gold-leaf, and in the 
finely divided state, as obtained by precipitation with ferrous sul¬ 
phate, for gilding porcelain. The gold is mixed with a suitable 
material and painted on the porcelain, and when the articles are 
afterwards heated in a kiln the metal burns into the glaze. Articles 
made of the base metals may be covered with a thin layer of gold 
(or gilt) in various ways. According to one method a gold amalgam 
is rubbed on the surface to be gilt, the mercury afterwards expelled 
by heating, and the rough surface polished by burnishing. In the 
wet way, a thin coating of gold is obtained by simple immersion in 
an alkaline solution of auric chloride, or better by a separate 
electric current. In the process of electro-gilding the object to be 
gilt is dipped in a warm solution of one of the double cyanides of 
gold and potassium and made the negative pole of an electric 
current. The coating of gold obtained in this way is far more 
durable and compact than that produced by any other process. 

The gold which is used by the goldsmith and for coinage is an 
alloy of the metal with copper and sometimes silver. This alloy 
not only melts more easily, but is also harder and does not so 
readily wear away when used. In England the fineness of a sample 
of gold is expressed in carats ; pure gold being said to be 24 carats. 
Thus 18 carat gold consists of 18 parts of gold alloyed with 6 parts 
of copper. English sovereigns and half-sovereigns are made of 22 
carat gold, or contain 916-66 parts of pure gold per thousand (or 
mille ). The gold used for coinage in Germany and some other 
countries has a fineness of 900, or contains 900 parts of pure gold 
per thousand, which corresponds to 21-6 carat gold ; 15 or 12 
carat gold is usually employed for the manufacture of jewellery. 


557 


PLATINUM METALS. 


This group of the noble metals includes, besides platinum, the 
five metals : iridium, palladium, rhodium, ruthenium, and osmium. 
They are almost always associated with one another in nature and 
are found native accompanying platinum. The so-called platinum 
ore contains them as alloys together with varying quantities of gold, 
copper, and iron. This ore is found in the Ural, in South America, 
California, Borneo, Australia, and a few other localities, and, like 
gold, usually in alluvial sands. 

These six metals closely resemble one another in many points, 
but in some respects differ strikingly. Such a difference is noticed 
for example, when we compare together their specific gravities and 
their atomic weights. It is then seen that the three: platinum, 
iridium, and osmium, possess a specific gravity and atomic weight 
which are about double those of the other three : palladium, 
rhodium, and ruthenium. These differences are exhibited in the 
following table:— 

Specific 
gravity 

• 21*5 
, . 22*4 


Platinum 
Iridium . 
Osmium 


22*5 


Atomic 

weight 

i 95 

i 93 

199 


Palladium 

Rhodium 

Ruthenium 


Specific 

gravity 

. n *4 

. I2‘l 
. I2'3 


Atomic 

weight 

106 

104 

104 


A second difference between these metals is their behaviour 
with regard to oxygen. Most of them have so weak an affinity for 
Ixygen that they not only do not unite directly with this element, but 
the majority of their oxides, when prepared in an indirect manner 
lose *eir oxygen when heated. On the other hand, osmium and 
ruthenium combine energetically with oxygen when simply heated 
in the air and produce volatile compounds of an acid character. 

Of the six rare platinum metals, platinum itself is found in far 
larger quantities than the other five. On this account, and because 
the metal is used for many purposes in the arts, the compounds of 
platinum are better known than those of the other metals. But 
the majority of these compounds are only of chemical interest. 



553 


Text-Book of Inorganic Chemistry. 


PLATINUM. 

Chemical Symbol : Pt .—Atomic Weight : 195* 

In the extraction of platinum the ore is first warmed with 
dilute aqua regia, to dissolve any gold, and then treated with 
stronger aqua regia, often under pressure, when platinum with traces 
of iridium and rhodium go into solution as chlorides. The residue 
consists chiefly of the earthy impurities and of an alloy of osmium 
and iridium (osmiridium) which is not attacked by aqua regia. 
The solution is evaporated to dryness, then digested with water, 
and the clear liquid mixed with a solution of ammonium chloiide. 
A yellow crystalline double chloride of platinum and ammonium 
is thus produced which is difficultly soluble in water and insoluble 
in alcohol. The salt is filtered off, washed, dried, and then heated 
to redness, when metallic platinum remains behind as a grey spongy 
mass, called spongy-platinum. This is the material used for the 
preparation of vessels made of platinum j it still contains traces 
of iridium (and rhodium). Very various and often complicated 
processes are used to completely separate platinum from all the 
other platinum metals and to prepare it perfectly pure, for a de¬ 
scription of which the student must consult a larger book. Small 
quantities of pure platinum may be prepared by the following 
method, due to Bunsen. The solution of the chlorides is mixed 
with a large excess of caustic soda until the precipitate which is first 
produced is redissolved and then boiled with the addition of a few 
drops of alcohol. During this boiling, the platinum tetrachloride 
remains unchanged, but the tetrachlorides of the other metals arc 
reduced to a dichloride or a sesqui-chloride, and these do not yield 
an insoluble compound with ammonium chloride. When, therefore, 
the alkaline solution is afterwards acidulated with hydrochloric 
acid and mixed with an excess of ammonium chloride, the platinum 
alone is precipitated, and the other metals remain in solution. 

Platinum is a greyish-white metal of specific gravity 21-5 ; it is 
malleable and ductile, and very tough, and indeed, scarcely less so 
than iron. Small quantities of iridium render it less malleable, 
but at the same time make it harder and tougher and more suitable 
for the manufacture of crucibles, dishes, &c. It cannot be fused 
at the temperature of our furnaces, but readily melts in the oxy- 
hydrogen flame, when heated in a lime crucible (p. 24). Thin 


Platinum. 559 

platinum wire melts easily in the oxy-hydrogen blowpipe, and is 
even partly volatilized. 

Platinum shares with iron the property of welding at a high . 
temperature. Two pieces of platinum heated to whiteness, but 
still far below the melting point of the metal, may be welded 
together under the hammer. In the same way, the spongy platinum 
which remains behind when ammonium-platinum chloride is heated, 
may be welded into a compact mass under strong pressure at a 
high temperature. Such platinum, however, still contains traces 
of air, and crucibles made of it often become covered with small 
excrescences due to the expansion of this confined air. 

Platinum may be obtained in a still more finely divided state 
than it exists in spongy platinum. We are acquainted with the 
metal in the form of a loose very fine black powder, called pla¬ 
tinum black. This form of platinum may be obtained by the re¬ 
duction of an alkaline solution of platinic, or better, platinous 
chloride. For this purpose platinous chloride is dissolved in an 
excess of caustic potash (or soda), the mixture heated to boiling, 
and small quantities of alcohol gradually added. The metal is 
then reduced and the alcohol oxidized to aldehyde and acetic acid. 
This black powder, when carefully washed and dried at a low 
temperature, absorbs and condenses considerable quantities of 
various gases, especially oxygen, and is therefore a powerful oxi¬ 
dizing agent. 

Spongy platinum also possesses this property of condensing gases, 
but in a less degree. If a jet of hydrogen is allowed to play on a 
small piece of spongy platinum in the air, the platinum first becomes 
red hot and then ignites the hydrogen. Dobereiner’s hydrogen 
lamp depends upon this property of spongy platinum. The oxygen 
which is condensed in the spongy platinum has stronger affinities 
than the ordinary gas, and as platinum also condenses hydrogen, 
the two substances are brought close enough together to cause their 
union. 

Compact metallic platinum offers a far less surface to the gases 
than the spongy metal; it acts therefore less energetically, and not 
at all at the ordinary temperature. But if a spiral of platinum 
wire or a piece of thin platinum foil is heated to low redness and 
suspended in a vessel containing a little ether, the vapour of the 
ether, which is contained in the vessel mixed with air, gradually 
becomes oxidized to aldehyde and acetic acid, with so much 
evolution of heat that the platinum spiral remains glowing, and 


560 Text-Book of hiorganic Chemistry . 

that the ether sometimes catches fire. It has already been men¬ 
tioned (p. 162) that compact platinum, when heated, can convert a 
mixture of sulphurous anhydride and oxygen into sulphuric anhy¬ 
dride without itself undergoing any change. Platinum black, con¬ 
taining the metal in its most finely divided state, acts more power¬ 
fully than either the spongy or compact metal. It can absorb as 
much as 200 times its volume of oxygen, or even more, and we 
may therefore suppose that the gas is condensed by it to the liquid 
or even solid state. If platinum black is moistened with alcohol 
and exposed to the air, the latter substance becomes oxidized to 
aldehyde and acetic acid, and so much heat is set free that the 
metal is raised to redness and continues to glow for some time. 

Platinum is not attacked by either nitric, hydrochloric, or sul¬ 
phuric acid, but dissolves in aqua regia with formation of platinic 
chloride. Molten caustic alkalies, especially lithia, attack platinum 
at a red heat, and a mixture of fused caustic potash and nitre acts 
more powerfully. Sodium or potassium carbonate is entirely with¬ 
out action on the metal even at a white heat. Sulphur, phosphorus, 
and arsenic also slowly attack compact platinum, and the spongy 
metal readily melts when heated with these substances. A mixture 
of silicates and charcoal also attacks platinum, with formation of 
a compound of platinum and silicon, and platinum crucibles are 
therefore soon destroyed if brought into direct contact with glowing 
charcoal. 

Many of the easily fusible heavy metals, such as lead, form 
easily fusible alloys with platinum. If, for example, a piece of lead 
is melted in a platinum crucible, an easily fusible alloy of platinum 
and lead is formed, and a hole is produced in the bottom of the 
crucible. Care must, therefore, be taken that these substances (sul¬ 
phur, phosphorus, arsenic, and the easily fusible heavy metals), or 
mixtures which might produce them, are never strongly heated in 
platinum vessels. 

The properties which are possessed by platinum—its infusibility 
at any temperature of our furnaces, that it does not unite with 
oxygen at any temperature, and that it resists the action of all acids 
including hydrofluoric acid, but excepting aqua regia—make the 
metal simply indispensable to the chemist. We employ it in the 
form of wire, foil, crucibles and dishes, tongs, tubes, retorts, &c. 
Considerable quantities of platinum are now used in the arts for 
the manufacture of large retorts and coolers in the purification of 
sulphuric acid. Platinum coins were also formerly used in Russia, 


Compounds of Platinum. 561 

but have now been withdrawn from circulation. Notwithstanding 
its scarcity, platinum is about one-third cheaper than gold, but is 
more than seven times as dear as silver. 


Compounds of Platinum. 

Platinum unites with other elements and produces two series of 
compounds—the platinous and platinic compounds—in which the 
metal is a dyad and a tetrad respectively. The compounds of 
platinum thus resemble those of carbon and tin. The most 
interesting of these substances are the chlorides and the peculiar 
platinum bases. 

Platinic Chloride : PtCl 4 .—On dissolving platinum in aqua 
regia and evaporating the solution to dryness with an excess of 
hydrochloric acid, a brownish-red crystalline and deliquescent mass 
remains behind of the composition: 2HCl,PtCl 4 or H 2 PtCl 6 + 
6 H 2 0 . If, however, excess of the acid is avoided, platinic chloride : 
PtCl 4 + 5 H 2 0 , is obtained in fine red crystals. Both dissolve in 
water and produce a reddish-yellow liquid with an acid reaction. 1 
Alcohol and ether also dissolve platinic chloride in the cold, but 
are oxidized with reduction of the chloride when the solution is 
warmed. When carefully heated platinic chloride loses one-half 
of its chlorine and becomes converted into platinous chloride ; if 
more strongly heated, it is completely converted into platinum and 
chlorine. 

Platinic chloride readily combines with the chlorides of the 
alkali-metals, of thallium and of the metals of the alkaline earths, 
producing double chlorides in which the chlorine combined with 
the platinum is twice as much as that united with the other 
metals. These salts are mostly difficultly soluble in water and 
insoluble in alcohol, and separate or crystallize out when solutions 
of the two chlorides are mixed. They usually crystallize in regular 
octahedra. 

These double chlorides have all a similar composition to the 
hydrochloric acid compound referred to above: 2HCl,PtCl 4 or 
H 2 PtCl 6 , and may be considered as salts of this chlorplatinic acid. 
Potassiu?n chlorplatinate : 2KCl,PtCl 4 or K 2 PtCl 6 , is a pale yellow 
crystalline precipitate produced when solutions of potassium and 

1 The brownish colour which a solution of platinic chloride sometimes 
possesses is due to the presence of traces of iridium and other impurities. —Ed. 

O O 





562 Text-Book of Inorganic Chemistry . 

platinic chlorides are mixed, and which is only completely precipi¬ 
tated on the addition of alcohol. When heated it breaks up into 
platinum chlorine and potassium chloride, the last of which may be 
extracted by boiling the residue with water. Ammonium chlorplati- 
nate\ 2NH 4 Cl,PtCl 4 or (NH 4 ) 2 PtCl 6 , closely resembles the potassium 
salt. When gently heated it leaves a residue of pure platinum as 
a grey spongy mass. Sodium chlorplatinate : 2NaCl,PtCl 4 or 
Na>PtCl 6 + 6 H 2 0 , is the only one of the double salts which crys¬ 
tallizes with water, and which dissolves both in water and alcohol. 
These properties enable us to completely separate sodium com¬ 
pounds from those of potassium. The sodium and potassium salts 
are first converted into chlorides, an excess of platinic chloride 
then added to this solution and evaporated to dryness. The 
residue is then extracted with strong alcohol when the sodium 
chlorplatinate and the excess of platinic chloride dissolve, leaving 
the insoluble potassium salt behind. It an excess of platinic chloride 
had not been added, some sodium chloride, which is insoluble in 
alcohol, would remain with the potassium chlorplatinate. Sodium 
chlorplatinate crystallizes from its concentrated aqueous solution 
in reddish-yellow prisms. 

Platinous chloride : PtCl 2 —If platinic chloride is carefully 
heated to 230° in a porcelain basin with continued stirring until 
no further odour of chlorine is perceived, it is completely converted 
into platinous chloride and chlorine. This latter compound is an 
olive-green powder, insoluble in water, but dissolving in hydro¬ 
chloric acid to form a dark-brown solution. It is converted by 
aqua regia into platinic chloride. Platinous, like platinic chloride, 
unites with other chlorides (of potassium, ammonium, &c.) and 
produces crystalline double salts, which are soluble in water. Its 
behaviour with ammonia shall be described under the account of 
the platinum bases. 

Platinic Oxide : Pt 0 2 , and Platinic Hydrate : Pt(OH) 4 .—If 
a solution of platinic chloride is mixed with one of sodium car¬ 
bonate, the precipitate which is produced is not like platinic 
carbonate but a compound. of soda and platinic oxide. If the 
whole is then evaporated to dryness and extracted with dilute 
acetic acid, platinic hydrate remains behind as a compact reddish- 
brown powder. The compound does not dissolve in this acid, but 
produces soluble platinic salts with the stronger oxy-acids, which 


Compounds of Platinum. 563 

have as yet been but little investigated. The hydrate also unites 
with strong bases and produces compounds ( platinates ), in which 
it plays the part of an acid. When gently heated the hydrate is 
converted into a black powder of piatinic oxide , which is insoluble 
in acids. 

Platinous Oxide : PtO, and Platinous Hydrate : Pt(OH).,.— 
Platinous chloride when digested with caustic potash is converted 
into a dark-brown powder of platinous hydrate, which is insoluble 
in water, but dissolves both in strong caustic potash and in acids, 
producing brown solutions in both cases. When the hydrate is 
gently heated, platinous oxide remains as a black powder insoluble 
in acids. 

Piatinic Sulphide : PtS 2 , and Platinous Sulphide : PtS, are black 
precipitates produced when sulphuretted hydrogen is led through 
a solution of piatinic or platinous chloride respectively. They 
both behave as sulpho-acids and dissolve, for example, in potassium 
sulphide to form red solutions. 

Platinous Cyanide : PtCy 2 .—If platinous chloride is digested 
with sufficient potassium cyanide to dissolve it and the solution 
evaporated down, crystals of potassium-platinous cyanide : 2KCy, 
PtCy 2 or K 2 PtCy 4 , separate out. The crystals show a beautiful 
dichroi'sm— i.e. they appear of different colours when viewed in 
different directions. From a solution of this salt, dilute acids 
separate platinous cyanide as a greenish-yellow powder insoluble 
in water and acids. It unites with other cyanides and forms other 
double salts, similar in composition to the potassium compound, 
and which are partly soluble and partly insoluble in water. 


Ammoniacal Compounds of Platinum. 

Under this name is included a number of very interesting com¬ 
pounds, produced by the action of ammonia on platinous chloride 
and by further changes in the primary compound. In these com¬ 
pounds a portion of the hydrogen of ammonium is displaced by 
dyad platinum : Pt", or by the dyad radicals : (PtCl)", (PtO)", or 
even by ammonium itself. 

Substitution of the hydrogen in ammonia or in an ammonium 
compound by a dyad radical may take place in two different ways. 

002 


564 


Text-Book of Inorganic Chemistry. 


Either two atoms of hydrogen may be displaced in one molecule of 
ammonia by the dyad radical, or one atom of hydrogen may be 
symmetrically displaced from each of two molecules of ammonia 
by the radical. The compounds of the first kind are unstable, and 
readily pass into ordinary compounds of ammonia; such, for 

example, is cyanic acid (see p. 316). 

The second class of compounds—the diammonium compounds 
—are more stable and do not easily pass into compounds of 
ordinary ammonia. 

The ammoniacal compounds of copper, some of those of 
mercury, and possibly those of cobalt, belong to this class of diam¬ 
monium compounds, but far more important representatives of the 
same group are those which contain platinum. The simplest of 

these is : ptatinoso-diammonium chloride-. Pt " a 

compound consisting of two molecules of ammonium chloride in 
which two atoms of hydrogen (one from each molecule) are displaced 
by one atom of dyad platinum. And the following compounds 
may serve as types of this very numerous group . 


(i.) Platinoso - diammonium \ 
chloride[ 

(ii.) Diammon - platinoso -1 
diammonium chloride) 


Pt" 

Pt" 


NH.C 1 

NH3CI 

NH 2 (NH 4 )C 1 

NH 2 (NH 4 )C 1 


(iii.) Diammon - platinoso -1 
diammonium hydrate [ 


fNH 2 (NH 4 )OH 
Pt |nh 2 (nh 4 )OH 


(iv.) Chlorplatin-diammonium) 
chloride [ 


(v.) Diammon - chlorplatin- j 
diammonium chloride 1 


PtCl a " 

PtCl" 


NH3CI 

NHgCl 

NH.,(NH 4 )C 1 

NH.;(NH 4 )C 1 


(vi.) Oxyplatin - diammonium) 

nitrate} 

(vii.) Diammon - oxyplatin-j 
diammonium nitrate f 


. PtO" 
. PtO" 


(NH 3 - 0 N 0 2 

|nh 3 -ono 2 

fNH,(NH 4 )-ON 0 2 

{nh;(nh 4 ).ono 2 


These salts are partly crystalline and soluble in water. The 
following is a brief account of the methods used for obtaining 
them. 

If ammonia is gradually added to a solution of platinous cmo- 
ride in dilute hydrochloric acid, a green crystalline salt separates 
out, which is a double compound of diammon-platinoso-diammonium 
chloride with platinous chloride 



Ammoniacal Compounds of Platinum. 565 


Pt 


NH„(NH 4 )C1 

NH~(NH 4 )C1 


+ 


PtCl 2 . 


This is commonly known as the green salt of Magnus , because 
it was discovered by this chemist \ it may also be prepared from 
its separate constituents. If this double salt, or if platinous chlo¬ 
ride (dissolved in hydrochloric acid) is heated with an excess of 
ammonia, the compound diammon-platinoso-diammonium chloride: 

(ii.) : Pt// |NH'(NH 4 )C1 cr y stalbzes out * n bne neecb es. Tbe 
lowing equation illustrates its production :— 


PtCl 2 + 4 NH 3 


Pt" 


NH 0 (NH 4 )C1 

NH 2 (NH 4 )C1 


If the two atoms of chlorine which this salt contains are displaced 

by hydroxyl, the corresponding base (111.): Pt |nH 2 (NH 4 )OH 1S 

produced; this crystallizes from its solutions and caustic potash 
expels no ammonia from it. It is a powerful base, and forms 
numerous salts with acids. When this base & heated, it parts with 
2 molecules of ammonia and is converted into platinoso-diajnmonium 


hydrate : a substance which is no lon S er soluble in 

water, and which gives insoluble explosive compounds with acids. 
Similarly, if the chloride (ii.) is heated to a temperature not exceed¬ 
ing 270°, it also loses 2 molecules of ammonia, and becomes 

platinoso-diammonium chloride (i.) : PP'j^jHjCl a sab wb * cb 

s only slightly soluble in water, and which readily again takes up 
2 molecules of ammonia when boiled with this substance. 

If chlorine is led into water which contains platinoso-diammo¬ 
nium chloride in suspension, each molecule of the salt combines 
with 2 atoms of chlorine and the dyad platinum is changed into 
the dyad radical (PtCl 2 )", containing the metal in the tetrad form. 
The salt then produced is chlorplatin-diammonium chloride (iv.), 
according to the equation :— 


Pt" 


NHgCl 

NH 3 CI 


+ 


CL 


(Pt iv Cl 2 )" 


{NHgCl 

(NH 3 CI 


This salt, which thus contains its 4 atoms of chlorine differently 
united, is only slightly soluble in water, and when boiled with 
ammonia takes up 2 molecules and becomes diammon-chlorplatin- 


566 Text-Book of Inorganic Chemistry. 

diammonium chloride (v.) : (PtCl) 2 " a salt which 

may be also produced by leading chlorine into a solution of diam- 
mon-platinoso-diammonium chloride (ii.). 

Oxyplatin-diammoniumnitrate(vi.): (Pt^O)"] NH 3 - 0 N 0 2 ma ^ 

be obtained from chlorplatin-diammonium chloride (iv.) by long 
boiling with an excess of silver nitrate, and when heated with 
ammonia yields diammon-oxyplatin-diammonium nitrate (vii.) : 

(PtOr'(gg^H® cr y staIlizes we ^ f rom hot water. 

Similarly constituted diammonium compounds containing a 
dyad radical are known in organic chemistry. 


PALLADIUM. 

Chemical Symbol : Pd .—Atomic Weight : 106. 

Platinum ores only contain a very small quantity of palladium— 
at most one or two per cent., and the metal is usually prepared 
from a gold ore which contains it, and which occurs in Brazil. 
This ore is fused with silver, then granulated and digested with 
nitric acid, which extracts the silver and palladium and leaves the 
gold behind. The solution is mixed with sodium chloride to preci¬ 
pitate the silver and the palladium then thrown down from the filtrate 
as palladous cyanide by the addition of mercuric cyanide. The 
cyanide yields the metal when heated to redness, which may then 
be fused in the oxy-hydrogen flame. Palladium melts more easily 
than platinum, and is, in fact, the most easily fusible of all the plati¬ 
num metals, but can only be melted with considerable difficulty in 
a wind furnace. Palladium resembles platinum in its external 
appearance, but is only about half as heavy : its specific gravity is 
n*4- ^ Is further distinguished from platinum by its solubility in 

nitric acid, and it is even attacked by warm concentrated sulphuric 
acid. It is malleable and ductile, and may be welded at a white 
heat. When heated in the air the metal tarnishes, but regains its 
original lustre when the temperature is raised to a bright red heat. 

Palladium Hydride : Pd 3 H.—The most remarkable property 
of palladium, a/property which is not possessed by any other metal 



Palladium . 


567 


to a similar extent, is its power of absorbing and retaining large 
quantities of hydrogen. The metallic appearance of the compound 
of palladium and hydrogen reminds one of an alloy, and appears 
to prove that hydrogen is a gaseous metal. Its composition is 
Pd 2 H, which corresponds to 600 volumes of hydrogen absorbed by 
one volume of the metal. The compound is formed at the ordi¬ 
nary temperature when palladium is brought into an atmosphere 
of hydrogen, but the most favourable temperature for its formation 
is ioo°. Above 130° and under the ordinary atmospheric pressure 
it begins to decompose with evolution of hydrogen, but the gas can 
only be completely expelled by protracted glowing. Besides the 
hydrogen which is chemically united with the metal, it can also 
absorb the gas mechanically. If water is electrolysed and palla¬ 
dium, especially if in the spongy form, used for the negative pole, 
the metal can absorb nearly 1,000 times its volume of the gas. But 
it readily loses the excess of hydrogen which it then contains 
beyond that combined chemically, and at ioo° the whole of the 
hydrogen mechanically absorbed is expelled. If a piece of palla¬ 
dium which has thus been saturated with hydrogen is exposed to 
the air, the gas is rapidly oxidized to water, with a considerable 
evolution of heat; and if placed in a solution of potassium ferricy- 
anide the salt is rapidly reduced to potassium ferrocyanide. Such 
palladium hydride when saturated with hydrogen is, in fact, one of 
the most powerful of reducing agents. It may be preserved for a 
long time unchanged if quickly placed in cold boiled water or 
alcohol. 

This property which palladium possesses of combining both 
chemically and mechanically with hydrogen, explains why hydrogen 
when led through a palladium tube heated to 130° diffuses through 
the walls as if they were porous. At a red-heat hydrogen diffuses 
still more readily through the metal. 

Palladium plays the part of a dyad and tetrad in its compounds. 
We are acquainted with a palladous oxide : PdO, and a corre¬ 
sponding palladous chloride : PdCl 2 , with a palladic oxide : Pd 0 2 , 
and with palladic chloride : PdCl 4 . 

On dissolving the metal in nitric acid, a brown solution is pro¬ 
duced which contains Palladous nitrate'. (N 0 2 ) 2 0 2 Pd, and which 
yields yellowish brown deliquescent prisms when evaporated to a 
syrup. The solution serves to detect iodine and to separate this 
substance from bromine and chlorine, since palladous iodide is 
quite insoluble in water, while the corresponding bromide and 


568 Text-Book of Inorganic Chemistry. 

chloride dissolve easily. If the dry salt is carefully heated as long 
as nitrous fumes are evolved , palladous oxide : PdO, remains behind 
as a black powder scarcely soluble in acids. 

, If palladium is dissolved in aqua regia, a dark-brown liquid 
results, which contains palladic chloride : PdCl 4 ; and if this solution 
mixed with the alkaline chlorides, similar double salts to those 
containing platinic chloride are produced. 

Potassium Chlorpalladinate : 2KCl,PdCl 4 or K 2 PdCl 0 , crystal¬ 
lizes in dark red octahedra. When small quantities of caustic potash 
are added to this salt a dark yellow powder of palladic hydrate 
containing potash is produced. And this when washed with warm 
water leaves a residue of black palladic oxide : Pd 0 2 , which is 
scarcely attacked by oxy-acids. 

When the solution of palladium in aqua regia is evaporated 
down, it loses chlorine and is converted into palladous chloride , 
which crystallizes out when the solution is slowly evaporated as 
brownish-red prisms of the composition : PdCl 2 + 2H 2 0, and which 
do not deliquesce in the air. This chloride also produces crys¬ 
talline double salts with the alkaline chlorides. A solution of 
palladous chloride in water is a very delicate reagent for carbonic 
oxide. The smallest quantities of this gas when brought into 
contact with the solution cause reduction and separation of black 
palladium. 

Palladous Iodide : Pdl 2 , is quite insoluble in water and in 
dilute acids ; it is obtained as a black precipitate when a soluble 
iodide is added to a solution of palladous nitrate. The correspond¬ 
ing chloride and bromide are soluble in water. 

Palladous chloride unites with ammonia and produces com¬ 
pounds similar to the ammoniacal compounds of platinum (p. 563). 


IRIDIUM. 

Chemical Symbol : Ir .—Atomic Weight : 193. 

Iridium resembles platinum in many respects, and only differs 
from this metal in some of its less important properties. Its specific 
gravity is 22*4 ; and it is harder and more brittle, and much more 



Iridium. 


569 

difficultly fusible than platinum. In the compact state iridium is 
not acted upon by aqua regia, it is only dissolved when in a finely 
divided state as iridium-black. The metal may be oxidized by 
heating it strongly with a mixture of caustic potash and nitre. 

Iridium is contained in the so-called platinum residues which 
remain after the ore has been digested with aqua regia. It is 
usually present united with osmium as an alloy of osmiridium in 
the form of very hard granules. This alloy may be decomposed by 
mixing it with common salt and heating the mixture in a stream of 
moist chlorine. The osmic chloride which is then produced is 
decomposed by the water contained in the moist gas, and converted 
into osmic acid, which is a volatile compound and which condenses 
in a receiver attached to the other end of the tube in which the 
mixture is heated ; at the same time the iridium becomes sodium- 
iridium chloride. The residue remaining in the tube and con¬ 
taining this salt is extracted with water, mixed with excess of 
sodium carbonate and evaporated to dryness. The dry mass when 
heated to redness yields iridium sesquioxide, which remains when 
the sodium chloride and carbonate are dissolved in water, and 
which is easily reduced to metallic iridium when gently heated in a 
stream of hydrogen. 

Besides the compounds of iridium with oxygen and chlorine 
which correspond to those of platinum with the same elements 
i.e. iridic oxide : Ir 0 2 , indious oxide : IrO, iridic chloride . IrCl 4 , 
and iridouschloride : IrCl 2 , it also forms iridium sesquioxide : lr 2 0 3 , 
and iridium sesquichloride : Ir 2 Cl 6 . The former of these com¬ 
pounds remains as a black powder when sodium-indium chloride 
is fused with sodium carbonate, and the fused mass extracted with 
water (see above). Iridium sesquichloride is obtained when iridium 
is strongly heated in chlorine as a yellowish-brown amorphous 
mass, insoluble in water and acids. 

Iridic Chloride : IrCl 4 , is obtained by dissolving the sesquioxide 
in aqua regia. It is much darker in colour than platinic chloride, 
and its solution is of a dark-brown colour. It produces double salts 
with the alkaline chlorides similar to those of platinum, but much 
darker in colour. Potassium-iridium chloride : 2KCl,IrCl 4 or 
K IrCl forms small dark red, nearly black octahedra, insoluble 
in" cold’water and only dissolving slightly when warmed. The 
ammonium salt is very similar. 




570 


Text-Book of Inorganic Chemistry. 


RHODIUM. 

Chemical Symbol : Rh.— Atomic Weight : 104. 

This rare element only occurs in very small quantities in plati¬ 
num ores, and has therefore been but little studied. In its chemi¬ 
cal properties it is most closely related to iridium. Rhodium is a 
malleable and ductile metal, more easily fusible than iridium, but 
less so than platinum, and is insoluble in aqua regia, except when 
alloyed with platinum, in which case it dissolves with this metal. 
When heated in the air it oxidizes slightly. Its oxides and chlo¬ 
rides correspond to those of iridium ; the most stable are rhodium 
sesquioxide : Rh 2 0 3 , and rhodium sesquichloride : Rh 2 Cl tf . 


OSMIUM. 

Chemical Symbol : Os.— Atomic Weight: 199. 

The separation of this metal from iridium has been described 
on the preceding page. If the vapours of osmic acid are led into 
caustic potash, the resulting liquid evaporated to dryness with much 
ammonium chloride, and the dry residue heated to redness, metallic 
osmium remains behind as a black powder when the fused mass is 
extracted with water. The metal remains unmelted at the highest 
temperatures; even in the oxy-hydrogen flame it only becomes a 
little more dense. It may be obtained as a hard crystalline powder 
when fused with tin, and the tin afterwards removed by hydrochloric 
acid. Osmium has a specific gravity of 22-5, and is apparently the 
densest of all substances. Aqua regia, or even nitric acid, oxi¬ 
dizes osmium to osmic acid. In the finely divided state, the metal 
burns when heated in the air, and produces the same substance. 

The compounds of osmium with oxygen and chlorine mostly 
correspond to those of platinum, thus : osmous oxide : OsO, and 
osmic oxide : 0 s 0 2 , osmoits chloride : OsCl 2 , and osmic chloride : 
OsCl 4 , and besides these an osmous atihydride : 0 s 0 3 . But by far 
the most important and interesting of the osmium compounds is 

Osmic Anhydride or Osmic Acid : 0 s 0 4 , in which the metal 
behaves as an octad element. This compound is produced when 
osmium is heated in a stream of oxygen, and is then deposited in 



Ruthenium. 


571 


the cooler parts of the tube in colourless crystals. It is also ob¬ 
tained when the metal or its alloy, osmiridium, is heated in a stream 
of moist chlorine. The corresponding chloride, which is then pre¬ 
sumedly first produced, but which is not known in the free state, is 
at once decomposed into osmic acid by the moisture of the damp 
gas. It melts under ioo° to a colourless liquid, boils at a slightly 
higher temperature, and sublimes in colourless glistening needles. 
Osmic acid possesses a piercing odour, somewhat resembling 
chlorine, and powerfully attacks the lungs and eyes. Great care 
must therefore be exercised in experimenting with this substance. 
Osmic acid, or more correctly, osmic anhydride, dissolves in water, 
but a compound with this liquid, which would be the true osmic 
acid, has not yet been prepared. Most reducing substances sepa¬ 
rate metallic osmium from its solution, a change which also goes 
on when the liquid is exposed to the light. A dilute solution is 
used by biologists for hardening organic tissues.— Potassium osmate 
is produced when osmic acid vapours are led into caustic potash, 
and separates from its solution as a crystalline powder on the addi¬ 
tion of alcohol. 


RUTHENIUM. 

Chemical Symbol : Ru.— Atomic Weight : 104. 

This metal, which resembles iridium and osmium, is less diffi¬ 
cultly fusible than the latter, but more so than the former metal. 
It has a specific gravity of 12-3, is scarcely soluble in aqua regia 
and insoluble in other acids. When fused with caustic potash 
and nitre it is converted into soluble potassium ruthenate , and 
when chlorine is led into a solution of this salt the temperature 
becomes high enough to volatilize the oxide ruthenic acid : Ru 0 4 , 
which is then formed. This compound then condenses in a cooled 
receiver as a yellow crystalline mass. In this respect, ruthenium 
resembles osmium, and the two metals are, at present, the only 
elements known of which one atom can unite with four atoms of 
oxygen. 







■ 


' 







APPENDIX. 1 


o 


Atomic and Molecular Weights. 

The atomic weights of the elements play so important a part both in theo¬ 
retical and in practical chemistry, that a brief account of the methods used to 
obtain these numbers may not be out of place, even in an elementary text-book 
of the science. 

In deciding upon the number which shall be considered as the atomic weight 
of an element, two distinct questions must be experimentally answered. ' In the 
first place, the weight of the element must be found which unites with a given 
weight of some other element or elements, of which the atomic weight or 
weights are accurately known ; and, secondly, it must be decided whether the 
number, so obtained, is really the atomic weight, or whether it is only some 
simple multiple or fraction of it. It is hardly necessary to remind the student 
that the atomic weights are simply relative and not absolute numbers, and that, 
since the atomic weight of hydrogen is taken as unity, the atomic weights of the 
other elements simply express how much heavier their atoms are than those of 

^ Ttefirst thing to be done in determining the atomic weight of an element 
is therefore, to find the ratio (chemical equivalent) in which it combines with 
other elements. For this purpose, the element is either made to unite with 
other elements of which the atomic weights are accurately known, or else one 
of its compounds with these elements is decomposed and its composition accu¬ 
rately determined. Such experiments appear simple enough theoretically, but 
in practice are surrounded by many and often almost insuperable difficulties if 
accurate results are to be obtained. The element or its compound must be 
prepared in a state of absolute purity, and only perfectly pure chemicals must 
be used in all succeeding operations. The greatest care must be taken that no 
impurities are introduced either from the air or from the vessels employed, and, 
what is more difficult to avoid, that a constant error does not occur through the 
entire series of experiments and so vitiate them all. To eliminate this constant 
error as far as possible, it is advisable to use not one but several compounds of 
the element for these determinations, and to employ material which has been 
obtained from various localities. The chemist must have an accurate balance 


By the Editor. 







574 


Text-Book of Inorganic Chemistry. 

at his disposal, and must know how to use it properly ; he must find before¬ 
hand whether his set of weights are exact ratios of one another, for example, 
whether the io-gramme piece is exactly ten times as heavy as the i-gramme, 
and one-tenth as heavy as the ioo-gramme piece, and if they differ from this 
ratio, as the best weights always do, he must find how much they differ and 
make the necessary corrections in his weighings. Finally, he must add to all 
his results the weight of air displaced by the substance weighed, which will vary 
according to the specific gravity of the substance, and according to the height 
of the barometer, the amount of moisture contained in the air and the tempera¬ 
ture at the time of weighing. 

The above brief description will give some idea of the delicacy of the opera¬ 
tions which are necessary for a correct determination of the chemical equivalent 
of an element, and it may be safely affirmed that such investigations demand 
greater skill on the part of the operator than almost any other chemical experi¬ 
ments. 

Supposing, then, that the ratio in which the element combines with other 
elements (or its equivalent) has been accurately determined, it remains to find 
its true atomic weight. Let us imagine, for example, that a given weight of 
pure carbon (the diamond) was burnt in a stream of pure oxygen, that the car¬ 
bonic acid so produced was carefully collected and weighed by absorbing it in 
a series of potash-tubes, and that it was then found that 3'66 grammes of car¬ 
bonic acid were produced for every gramme of carbon burnt. It would then 
be known that 1 gramme of carbon, when burning to carbonic acid, combines 
with 3'66— 1 - oo = 2'66 grammes of oxygen, or that 6 grammes of carbon com¬ 
bine with 16 grammes of oxygen. If, .then, the atomic weight of oxygen 
is 16, that of carbon must be 6, or some simple multiple of this number. 
It would be 6 if the compound contained an equal number of atoms of 
each constituent, 12 if it contained twice as many oxygen atoms as those of 
carbon, and 3 if twice as many carbon as oxygen atoms were present. Or, 
again, suppose that silicon chloride were decomposed in some suitable manner 
so that its chlorine was converted into silver chloride, and that, knowing the 
atomic weights of silver and chlorine and the composition of silver chloride, it 
was found that 100 grammes of silicon chloride contain 83 ’53 grammes of 
chlorine. Such a result would show that 100 — 83'53 = 16^47 parts of silicon 
combine with 83'53 parts of chlorine, or that 7 parts of silicon unite with 
35'5 parts (one atom) of chlorine. Hence the atomic weight of silicon might 
be 7 or 14 or 28 or 3'5. In order to decide which of these numbers correctly 
represents the atomic weight, further experiments are necessary and other 
considerations must be brought to bear on the problem. The further con¬ 
siderations which assist us in obtaining a solution of this question may be 
classified under three heads, viz. :— 

1. The relations between the molecular weights of gases and their density 
(Avogadro’s law); 

2. The relations between the specific heats of the solid elements and their 
atomic weights (Dulong and Petit's law) ; and 

3. The relations between the composition of a solid compound and its crys¬ 
talline form (isomorphism). 


Appendix . 


575 


Let us new see how we can use these three laws to fix the atomic weights of 
the elements. 

I. The determination of the atomic weight of an element from the density of 
its gaseous compounds. 

The important law, first enunciated by Avogadro and therefore called after 
him, states that equal volumes of all gases whether simple or compound con¬ 
tain an equal number of molecules, or, in other words, that the density of a gas 
is proportional to its molecular weight (see p. 5°). Avogadro’s law is proved 
not only because it agrees with all known chemical facts, but also because it 
explains the physical properties of gases in the only possible manner. 

In the kinetic theory of gases, as worked out by Maxwell, Clausius, and 
others, it is supposed that a gas consists of a large number of small particles 
(molecules) moving at a high but variable velocity, and, therefore, constantly 
collidine with one another and with the walls of the vessel which contains the 
gas This collision of the molecules with the walls of the vessel produces the 
pressure which the gas exerts, or the pressure which must be applied to the gas 
to prevent it from expanding. If, now, we call m the mass of each molecule 
and v the the mean velocity with which the molecules are travelling, the aver¬ 
age kinetic energy possessed by each molecule is : \ my*. But if two sub¬ 
stances are at the same temperature there is no change in the temperature ot 
either when they are brought into contact with one another ; there is, it is true, 
an interchange of heat between them, but each gains as much heat as the other 
loses Hence if two average molecules of the same or different gases at the 
same temperature are brought together, no change in their temperature is pro¬ 
duced, and no change in their kinetic energy (measured as heat) takes place- 
i.e. the kinetic energy of each is equal, or 

% m x v 1 2 = h • ' • * M 

Again the total pressure which is exerted on the vessel containing the gas is 
represented bv the sum of the kinetic energy of all the molecules present in the 
vessel, and if we call this number n, and the total pressure P we have : 

p = \nm v 2 , 

or since we only measure the pressure if) in one of thethree directions in which 
it is exerted: ^ ....(*> 

And by experiment, the pressure exerted by equal volumes of all gases at 

the same temperature is equal, hence : 

p = p 2 , and \n x m x = \n^m 2 v 2 . 

But we have already seen (t) that at the same temperature, 

Therefore ~~ n<1 ' 

or the numler of molecules contained, in equal volumes of all gases at the same 

the p rcduc ‘ of ,h . e " umber a ° n f d 7 1 i- 

cute contained in unit volume into the mass of each-t.e. <*i-*.*i d *~ 

mm*. and since n 1 = n 2 , we get: 

d, *• d % = 


57 6 Text-Book of Inorganic Chemistry. 

or the density of a gas is proportional to the mass of its molecules or its molet ulai 
weight. 1 

This is a brief risumi of the physical proof of Avogadro’s law with which 
we started. 

It must now be decided with what standard the molecular weights are to be 
compared. We know by experiment that a given volume of oxygen, for 
example, weighs 16 times more than the same volume of hydrogen at the same 
temperature and under the same pressure, and, therefore, that the molecule 
of oxygen must be 16 times as heavy as the molecule of hydrogen, but we do 
not know how many atoms are contained in the molecule of each gas. That 
the molecule of hydrogen contains at least two atoms is clear from the fol¬ 
lowing. In the union of hydrogen and chlorine equal volumes of the 
two gases combine to form the same volume of hydrochloric acid. Or, one 
molecule of hydrogen unites with one molecule of chlorine and produces two 
molecules of hydrochloric acid. Each molecule of hydrogen (and of chlorine) 
is therefore broken up into two parts which cannot be less than the atom, 
but may contain more than one atom. Hence, therefore, the molecule of hy¬ 
drogen is considered to consist of 2 atoms; it is not impossible, but highl 
improbable, that its molecule may contain some multiple of this number of 
atoms. 

The atomic weight of hydrogen being taken as unity, and its molecule con¬ 
sisting of 2 atoms, its molecular weight becomes 2. And since air is 14*44 
times as heavy as hydrogen, the molecular weight of any gas is found by mul¬ 
tiplying its density (compared with air as unity) by 2 x 14*44 = 28*88 (p. 15). In 
this way the molecular weights of a number of the elements and of many com¬ 
pounds have been determined. But in making these determinations our ex¬ 
periments are confined to those substances which are volatile at comparatively 
low temperatures ; and, in the case of a compound, to those which do not 
completely decompose when heated above their boiling-point, or which do not 
dissociate or decompose when heated, and recombine again on cooling. And, 
from the experimental difficulties which surround the determination of many 
vapour densities as well as the deviation from Boyle’s and Charles’ law r s at 
temperatures only slightly above the boiling-point of a substance, it follows 
that the molecular weights obtained in this way are never so correct as those 
derived from the ordinary analysis of the substance. The density of the gas 
only tells us whether the molecular weight is a multiple or fraction of a certain 
number, and does not give us the exact number itself. We have seen (p. 574) 
that, by chemical analysis, silicon chloride contains 7 parts of silicon united 
with 35*5 parts of chlorine, and its molecular weight is therefore either 42*5 or 
some exact multiple of this. The density of this gaseous compound has been 
found to be 5*94, which, wffien multiplied by 28*88, gives 171*5, consequently its 
true molecular weight is 42*5 x 4 = 170*0, not 171*5, nor any other multiple of 
4 2 ' 5 - 

The following table gives the vapour densities (d) of those of the elements 
for wffiich this has been determined ; the third column contains the molecular 


1 Maxwell, Theory of Heat, pp. 289 et seq. 



Appendix. 


5 77 


weight (28-88 d) to which this points ; and the fourth the true molecular weight 
(M):- 


Element 

d 

28-88 y.d 

Hydrogen. 

0*0693 

2'00 

Oxygen. 

1-1056 

3 I- 93 

Nitrogen ...... 

0*9713 

28*05 

Chlorine. 

2-45 

7076 

Bromine. 

5‘54 

i6o*oo 

Iodine. 

872 

251-8 

Selenium (at 1,400°) .... 

5-68 

164-0 

Sulphur (at i,ooo°) .... 

2 24 

647 

,, (at 500°) .... 

678 

195 8 

Tellurium. 

9-08 

262-2 

Mercury ...... 

6-99 

201 "8 

Cadmium 

3'94 

113-8 

Phosphorus. 

4-48 

129-4 

Arsenic . ... 

10-42 

300-9 


M 


H* 

o 2 

n 2 

Cl* 

Br. 2 

h 

Se<> 

S 2 

S 6 

Te 2 

Hg 

Cd 

P 4 

As 4 


2 

3 2 

28 

71 

160 

254 

158 

64 

192 

256 

200 

112 

124 

3 °° 


From this table it will be seen that the majority of the elements contain two 
atoms in the molecule. To this rule, the hexatomic sulphur, at temperatures 
just above its boiling point, forms an exception, as well as the four elements— 
mercury, cadmium, phosphorus, and arsenic. Mercury and cadmium contain 
only one atom in their molecules, an assumption which is strongly supported 
in the case of mercury by certain physical properties of its gas; while the mole¬ 
cules of phosphorus and arsenic are made up of four atoms each. 

We possess, then, in Avogadro's law a valuable and certain means for find¬ 
ing the molecular weights of substances which can be converted into the 
gaseous state without decomposition ; but it remains to be seen how we can 
utilize the same generalization for determining the atomic weight of an element. 
In the first place, it is evident from the preceding table that the molecular 
weight of an element has no definite relation to its atomic weight, and we are 
not therefore justified in making any deduction from the molecular weight of 
an element as to its atomic weight. But the true atomic weight of an element 
may often be found from the molecular weights of its compounds. The atom 
of an element may be defined as the smallest quantity which can exist in 
any compound, and among a large number of compounds of one element 
it is extremely probable that the molecule of one or more of them will 
contain only one atom. If, therefore, we have determined the vapour densities, 
and consequently the molecular weights of a number of compounds of one 
particular element, we may take the atomic weight as the smallest quantity of 
the element which occurs in the molecule of any of these compounds. The 
result so obtained is never absolutely certain, because a compound might be 
discovered containing a less quantity of the element than its received atomic 
weight and it has but little value when the number of compounds compared 
together is small, but approaches certainty when the densities of a large number 
of compounds are known. The number obtained from this comparison is the 
maximum value possible for the atomic weight-it may be smaller, but it can¬ 
not be larger. p 




















578 Text-Book of Inorganic Chemistry. 

A few examples may serve to explain more clearly how the atomic weight of 
an element may thus be found. Among those elements of which numerous 
gaseous compounds are known is chlorine. If we examine all these com¬ 
pounds, we find that none of them contains less than 35’5 parts of chlorine in 
its molecule, and we therefore consider this number as the atomic weight of 
the element. 

The following table gives the percentage composition of four compounds of 
chlorine, their vapour density (d), the molecular weight indicated by this density 
(28‘88 x d ), and the true molecular composition (M) :— 


Compound 

Percentage 

composition 

d 

28-88 xd 

M 

Hydroch oric acid 

Mercuric chloride 

Phosphorous chloride 

Silicon chloride 

(H = 2741 

1 Cl = 97-26 f 
! Hg =73-801 

1 Cl = 26 '20 1 
j P = 22 561 
•Cl = 77 ' 44 ^ 
(Si =16-47) 

1 Cl = 83 73 ) 

1-247 

9-80 

4-88 

5'94 

36-0 

283-0 

140-9 

171-6 

iH = 1 "o 
•Cl =35-5 
f Hg = 200 
(Cl =71 = 35^x2 
(P =31 

l Cl =106-5 = 35^x3 
f Si =28 

1 Cl =142 = 35^x4 


The table shows that the chlorine contained in the molecules of the four 
compounds is in the proportion of 112:3:4, and as on examining a large 
number of other compounds of the element none has been found with less 
chlorine in its molecule than hydrochloric acid, we consider that this compound 
contains one atom of chlorine, mercuric chloride two atoms, &c. At present, 
therefore, the atomic weight of chlorine is taken as 35'5, but if, at any future 
time, a compound of chlorine shou’d be discovered with only i7'75 parts of 
chlorine in its molecule, the present atomic weight would have to be reduced 
to this number. In the same way, from the comparison of a still larger num¬ 
ber of vapour densities, the atomic weight of carbon has been fixed at 12, and 
that of oxygen at 16. And the corresponding numbers for many other elements 
have been determined with more or less certainty, according to the number of 
compounds examined. 

The atomic weights obtained from Avogadro’s law are increased in value 
by their agreement with those resulting from other considerations. They agree 
with those derived from the specific heats of the elements in the solid state, and 
from the general properties of the element and its compounds. 

II. The determination of the atomic weight of an element from its specific 
heat in the solid state. 

The specific heat of a substance is the quantity of heat required to raise 
1 gramme of the substance i c C. as compared with the quantity of heat required 
to raise 1 gramme of water through the same increment of temperature. The 
specific heats (or capacities for heat) of the various elements differ very con¬ 
siderably from one another. But if, instead of comparing together the quantities 
of heat required to raise equal weights of the elements i°, we compare those 
necessary to similarly increase the temperature of the atomic weights of the 
elements (in the solid state), we arrive at a different result: we then find that all 
these quantities of heat are nearly equal to one another. This important law, to 














Appendix. 579 

which we have already referred (p. 70), was discovered by Dulong&ndi Petit, and is 
therefore usually named after them. The law may also be expressed by saying 
that the capacity for heat of all the elementary atoms in the solid state {atomic 
heat ) is the same, or, that the product of the specific heat of a solid element 
and its atomic weight is a constant. The following table gives the specific 
heats of some of the elements in the solid state (s), the temperature {t) at which 
these quantities have been determined, the atomic weights {a. w.) of these 
elements and their atomic heats {a. h.) ; or the product of the specific heat and 
atomic weight in each case :— 


Element 

.s 

t 

a . w. 

a. h. 

Lithium. 

0*941 

0 

64 

7*0 

6*6 

Boron. 

0*366 

233 

11*0 

4*0 

Carbon (Diamond) 

0-459 

9 8 5 

12*0 

5 *5 

Silicon ...... 

0*203 

232 

28*0 

57 

Phosphorus. 

0*189 

19 

31*0 

5 '9 

Sulphur ..... 

0*171 

50 

32*0 

5'5 

Potassium. 

0*166 

-34 

39 *o 

6-5 

Iron. 

0*112 

+ 3 i 

56*0 

6 ‘3 

Copper. 

0*095 

58 

63*4 

6*o 

Silver. 

0*056 

50 

108 *o 

60 

Tin. 

0*056 

50 

u8*o 

6*6 

Platinum ..... 

0*0324 

55 

198*0 

6*4 

Lead. 

0 0315 

| 

55 

207*0 

6-5 


This table shows us, in the first place, that the atomic heats of the very 
various elements which it includes are all approximately equal, but we also see 
that the numbers differ considerably among themselves. These differences are 
principally due to the change of specific heat with temperature, which nearly 
always increases as the temperature rises. The most striking examples of this 
are seen in the elements carbon, boron, and silicon. At the ordinary tempera¬ 
tures the specific heats of these three substances are 
Carbon = 0*132 (at 33 0 ) 

• Boron = 0-238 (at 26°) 

Silicon = 0*170 (at 22 0 ) 

numbers which do not at all agree with the law— e.g. the atomic heat of carbon 
is then 0*132 x 12= 1*58. But as the temperature rises the specific heat of 
each rapidly increases until it becomes, to a certain extent, constant, and the 
atomic heat is then found to be normal. Thus, at high temperatures, the three 
elements have the following specific heats : 

Carbon = 0*459 ( at 9 8 5 °) 

Boron = 0*366 (at 233°) 

Silicon = 0*203 (at 232 0 ) 

For carbon and silicon these numbers correspond to the normal atomic 
heat, and that for boron probably would do so at a slightly higher tempe¬ 
rature. 

Dulong and Petit’s law is therefore only approximately true, or the specific 

P P 2 





















580 


Text-Book of Inorganic Chemistry. 


heat is a function not only of the atomic weight, but also of other properties of 
solids with which we are at present but little acquainted. But notwithstanding 
this, the law enables us to decide between one of two or more numbers, differing 
considerably from one another, in determining the atomic weight of an element. 
All that is necessary to find the required atomic weight is to divide the mean 
atomic heat (6-4) by the specific heat of the element in the solid state. The 
quotient is then the approximate atomic weight, which has to be corrected by 
ordinary analytical results. Two examples may be given of how the specific 
heat of a solid element controls its atomic weight. When the metal indium 
was discovered by spectrum analysis it was found in association with zinc, a 
metal which it closely resembles both in the free state and in its compounds. 
The only oxide which indium forms was therefore thought to correspond to 
the only oxide of zinc— i.e. to be a monoxide of the formula : InO ; and this 
gave an atomic weight of 75*6. But soon afterwards the specific heat of the 
metal was determined and was found to be 0-057, which gives an approximate 

atomic weight of 


^ 4 =xi2, and hence its true atomic weight is, not 75'6, 


°‘°57 

but 75’6 x 1 *5 = 113-4. From this it also follows that the formula of its oxide 
becomes ln 2 0 3 , and that the composition of all its other compounds must be 
correspondingly changed. Two atoms of indium now correspond to three of 
what were formerly considered as the atoms of the element. A precisely- 
analogous case is to be found in the atomic weight of the metal beryllium, 
which, during late years, has been taken as 9-1, and its oxide considered as a 
monoxide with the formula : BeO. Recently, however, the specific heat of the 
metal has been determined and found to be 0-445, which points to an atomic 

weight of 6 4 ■ = 14. The true atomic weight thus becomes q-ixri:=ir6; 

0-445 000. 

and the oxide is a sesquioxide : Be 2 0 3 . 


III. The determination of the atomic weight of an element from the crys¬ 
talline form of its solid compounds. 

It was discovered by Mitscherlich that many substances which crystallize in 
the same form with the same or nearly the same angles have a similar chemical 
composition. Such substances are said to be isomorphous with one another 
and this law of isomorphism may be used to determine the atomic weight of 
an element in the following way. If two substances crystallize in the same 
form and the composition of one is known, the composition of the other mav 
be inferred to be similar, and thus the atomic weight of any particular element 
which it contains may be fixed. The compounds alumina, ferric oxide, and 
chromic oxide all crystallize in the same form of the hexagonal system, and 
from the composition of one of them we can therefore deduce that of the other 
two : if ferric oxide = Fe 2 0 3> then alumina = A 1 2 0 3 , and chromic oxide = Cr 2 0 3 ; 
and from this ratio the atomic weights of aluminium and chromium may be 
deduced if that of iron is known. 

The atomic weights obtained by isomorphism are not, however, so trust¬ 
worthy as those based upon Avogadro’s or Dulong and Petit’s law, and for the 
following reasons. That the form in which a substance crystallizes does not 




Appendix. 5 ^i 


entirely depend upon its chemical composition is at once proved by the fact 
that the same compound may be dimorphous , Or may crystallize in two distinct 
and entirely different forms— e.g. calcium carbonate, arsenious anhydride, sul¬ 
phur, &c. In the next place, the composition of two compounds may be 
exactly similar, but they need not crystallize in the same form— e.g. the nitrates 
and the corresponding metaphosphates. And, finally, two substances may crys¬ 
tallize in precisely the same form without having a similar composition e.g. 
sodium nitrate (NaN 0 3 ) and calcite (CaC 0 3 ) ; sodium sulphate (Na 2 SC>4), 
and barium permanganate (BaM 2 03) ; cadmium sulphate (3CdSC>4+ 8 H 2 0 ), 
and didymium sulphate (Di 2 3S0 4 + 8 H 2 0 ). The evidence derived from isomor¬ 
phism is, therefore, only of value so far as it confirms the results derived from 
the other two laws, or when these two laws cannot be applied. Thus the 
isomorphism of the compounds :— 

K 2 SiF 6 + H 2 0 

K 2 TiF 6 + H 2 0 

K 2 ZrF 6 + H 2 0 

K 2 SnF 6 + H 2 0 


afforded valuable evidence concerning atomic weights of silicon, titanium, and 
zirconium, until they were confirmed by the other two more certain laws. 

The following is a list of some of the more important isomorphous groups, 
showing the chief elements which are connected together by this law: 


Groups. 

(i.) Phosphates 
Arsenates 

pi.) Carbonates of the composition :RC 0 3 , 
where R = Ca, Mg, Fe," Zn, Mn. 

(iii.) Double sulphates of the general 
formula : M' 2 S 0 4 + M"S04 + 6 H 2 0 , 
where 

M' =K, NH 4 

M"' = Fe", Co, Ni, Mg, Zn, Mn" 

(iv.) The Alums of the general formula: 
M'M"'3S0 4 + i2H 2 0, 
where 

M' = K,NH 4 , Rb, Cs, Na, Tl, Ag 
M'"=A 1 , Fe"', Mn'"Cr 
(v.) Chromates, Manganates, Sulphates, 
and Selenates 
(vi.) KMn 0 4 and KC 10 4 
(vii.) Ag 2 S 0 4 , Na 2 S 0 4 
(viii.) K 2 SiF 6 , H 2 0 ; K 2 TiFe, H 2 0 , 

K 2 ZrF 6 , H 2 0 ; K 2 SnF 6 ,H 2 0 . 

(ix.) Apatite : 3Ca 3 P 2 08, Ca(Cl,F) 2 
Pyromorphite : 3Pb 3 P 2 0 8 , PbCl 2 
Mimetesite : 3Pb 3 As 2 08. PbCLj 
Vanadinite : 3Pb 3 V 2 08iPbCl 2 


Elements Connected. 
P and As 


Ca, Mg, Zn, Fe, Mn 


Fe",Co,Ni, Zn, Mg, Mn'' 


K,(NH 4 ),Rb.Cs,Na.Tl, Ag 


Al, Fe'", Mn'", Cr" 
Cr vi , Mn", S, Se, 


Mn™, CP 11 
Ag, Na 
Si, Ti, Zr, Sn 

P, As, V 

Pb, Ca 


582 


Text-Book of Inorganic Chemistry. 


Relations between the Atomic Weights. 

It has been already remarked (pp. 65 et seq.) that certain of the elements 
which possess similar chemical properties and which usually occur associated 
together in nature may be arranged in groups of three ; and that when this is 
done the atomic weight of one is about the mean of the other two. The fol¬ 
lowing have been mentioned as examples of this triple grouping :— 


(i.) 

Chlorine 

= 

35’5) 



Bromine 

= 

80 -o - Mean 

= 81-25 


Iodine 


127-0) 

(ii.) 

Sulphur 

= 

3 2 '°) 



Selenium 

= 

79-0 J- Mean 

= 80'O 


Tellurium 

= 

128-0 j 


(iii.) 

Lithium 

= 

7 ’°j 



Sodium 

= 

23-0 r Mean 

= 23-0 


Potassium 

= 

39‘° i 


These relations cannot be accidental, and it certainly does appear that the 
properties of an element and of its compounds are related in some way or 
other to its atomic weight. 

Of the various numerical relations which have been supposed to exist be¬ 
tween the atomic weights we shall here only refer to one hypothesis, which was 
put forward by Prout, and is commonly known as Prout's law. With regard 
to the connexion between the atomic weight of an element and its properties, 
the most important theory is that known as the Periodic law. This theory, 
which has been developed in late years by Newlands, Lothar Meyer, and 
Mendelejeff, supposes that the properties of the elements change periodically as 
we pass from one to the other when arranged according to their atomic weights. 

Prout's Law .—The hypothesis which is known under this name supposes 
that the atomic weights are all multiples of that of hydrogen by whole numbers, 
and that the elements may therefore be considered as compounds of hydrogen 
of varying degrees of complexity. Prout’s hypothesis was advanced at a time 
(1815) when the atomic weights were much less accurately known than they 
are at present, and when they were usually represented by the nearest whole 
numbers to which the results pointed. It soon became clear, however, that the 
atomic weights of certain of the elements, notably chlorine, could not be re¬ 
presented by whole numbers, and in order to make the hypothesis applicable 
to these it was proposed to call the atomic weight of hydrogen 2, or even 4. 
But even with such devices as this, accurate determinations of the atomic 
weights of many of the elements by which every possible error was avoided or 
allowed for have shown that in most cases the hypothesis is certainly not true. 

The atomic weights which are given on p. 60, and used in the text of this 
book, have been mostly taken as whole numbers to avoid fractions ; but for 
some elements much more accurate results than this have been obtained by 
careful investigators and have been confirmed by others. The following table 


Atpendix. 5^3 

gives a list of these more accurate atomic weights, which shows that only in a 
few cases is Prout’s hypothesis approximately true. In the second column the 
ratio between the atomic weights of hydrogen and oxygen is i : 15-96, and in 
this column the atomic weights are calculated for these numbers. We may, 
however, consider the atomic weight of oxygen as 16, when all the other atomic 
weights will be slightly larger. These numbers are given in the third column. 
It will be noticed that the atomic weights there more nearly approach whole 
numbers, and hence Prout's law has not yet been completely forsaken by some 
chemists. 


Table of Accurately Determined Atomic Weights. 


Element 

0 = 15*96 

0 = 16 

I Aluminium ....•• 

Antimony. 

Barium. 

Boron ...••• 

Bromine 

Cadmium. 

Caesium ..•••• 

Calcium. 

Carbon. 

Chlorine. 

Copper. 

Hydrogen ..•••• 
Indium ..•••• 

Iodine ..•••• 

Iridium. 

Iron. 

Lead. 

Lithium. 

Mercury. 

Nitrogen. 

Oxygen ..•••• 

Phosphorus. 

Platinum ..•••• 

Potassium. 

Rubidium. 

' Selenium. 

Silver. 

Sodium. 

Strontium. 

Sulphur. 

Vanadium. 

27-02 

119-8 

136-81 

10- 92 

7976 

11175 

132-65 

39'95 

11- 97 

35'37 

63-18 

I’OO 

H 3'4 

126-55 

i 9 2 '5 

55'9 

206-43 

7 ’ CI 

1997 

14- 01 

15- 96 

30-96 

I 94'5 

39-02 

85-2 

78*8 

107-67 

22-99 

87'3 

3 i ’98 

5 1 ’ 1 

27-09 

120*1 

I 37 'I 5 

10-95 

79-96 

112-03 

132-98 

4005 

12 -OO 

35*46 

63*34 

10025 

1137 

126-87 

192-98 

56-14 

206-95 

7*03 

200-8 

14*05 

16-00 

3 1 *°3 
194*99 

39-12 

85*4 

79 *o 

107*94 

23*05 

87*5 

32-06 

5 1 ' 2 


The Periodic Law. 

If the elements are arranged according to their atomic weights commencing 
with the smallest, it is found that those elements with similarchemicaland physical 
properties occur at regular intervals. Various tables have been drawn “P “ 
exhibit these periodic changes in the properties of the elements, one of which 





























5 §4 


Text-Book of Inorganic Chemistry. 

is given below, and the series of facts thus shown to be connected with the 
atomic weight are collectively known as the periodic law. We owe this useful 
generalization principally to the labours of Lothar Mejer and Mendelejeff:— 


I 

II 

III 

IV 

Y 

VI 

VII 

VIII 

H 










i 




1 ' 






Li 

Be 

B 

c 

N 

O 

F 




7 

? 

11 

12 

14 

16 

*9 




Na 

Mg 

A1 

Si 

P 

s 

Cl 




2 3 

24 

27 

28 

3* 

3 2 

35‘5 




K 

Ca 


Ti 

V 

C r 

Mn 

Fe 

Ni 

Co 

39 

40 


48 


52 

'55 

56 

58-5 

59 

[Cu] 

Zn 

Ga 


As 

Se 

Br 




6 3 

6 5 

70 


75 

79 

80 




Rb 

Sr 

Y 

Zr 

Nb 

Mo 


Ru 

Rh 

Pd 

85 

88 

? 

90 

94 

96 


103 

104 

106 

[ J g] 

Cd 

In 

Sn 

Sb 

Te 

I 




108 

112 

113 

118 

120 

? 

127 




Cs 

Ba 

La 

Ce 

Di 






i33 

i37 

138 

x 4 r 

*45 








Er 

Yb 

Ta 

W 


Os 

Ir 

Pt 



166 

w 

10 

182 

184 


? 

i93 

iQ 5 

[AuJ 

Hg 

T1 

Pb 

Bi 






199 

200 

204 

207 

210 









Th 


u 








231 


2^0 






Referring to the accompanying table, which commences with hydrogen and 
ends with uranium, we find that, omitting hydrogen, the first seven elements 
correspond in their properties to the next following seven : thus, Li corre¬ 
sponds to Na, C to Si, N to P, &c. Following on in the series we see again 
that the next seven (less one) also correspond in a regular manner to the pre¬ 
ceding seven. In this way the elements become classified into various groups, 
which are shewn in the vertical columns of the table, and which contain closely- 
related elements. Group I., for example, contains Li, Na, K, Rb, Cs, and the 
allied metals Ag and Cu, as a sub-group ; Group II. includes the metals of the 
alkaline earths, together with Mg, Zn, Cd, and Hg. The elements in each of 
these groups are not only allied chemically but also often connected by similar 















































































































Appendix. 5^5 

physical properties ; they further possess the same atomicity in some of their 
compounds. The elements included in the first group are monads, those in the 
second group are dyads, and so on. 

This brief description of the periodic law, which space does not allow us to 
expand more fully, shows clearly that the properties possessed by an element are 
closely connected with its atomic weight ; and that the method in which the 
elements are arranged is to some extent correct, is corroborated by other con¬ 
siderations. It will be noticed that there are several gaps in the above table— 
e.g. between Ca and Ti, Ga and As. Mo and Ru, neglecting those among the 
elements of higher atomic weights. These gaps are supposed to be occupied 
by undiscovered elements, a supposition which received a remarkable confir¬ 
mation in the recent discovery of the metal gallium. Mendelejeff had pre¬ 
viously predicted that such a metal might exist, and had described its properties 
and approximate atomic weight under the name of ekaluminium. And when 
a few years afterwards the metal now called gallium was discovered, its pro¬ 
perties were found to agree almost exactly with Mendelejeff s hypothetical ek¬ 
aluminium. The periodic law has also done good service in correcting some 
uncertain atomic weights. Up to a recent date the metals of the cerium and 
yttrium groups were considered as dyads and their ox'des as monoxides , but 
Mendelejeff showed that the atomic weights which they then possessed did not 
agree with their properties according to the periodic law, and proposed to 
multiply these numbers by i -5, so that the oxides would become sesquioxides. 
The atomic weights which he had proposed for the cerium metals were after¬ 
wards found to be correct as soon as their specific heats were determined. 
In the same way the atomic weight of uranium w r as formerly thought to be 
120, but was supposed by Mendelejeff to be 240—a supposition which was also 
confirmed by a determination of the specific heat of the metal. 

On the other hand, it is none the less true that the periodic law fails to find 
a position for some elements of which the atomic weight has been confirmed by 
Dulong and Petit’s law, or by careful analysis. According to the specific heat 
of beryllium its atomic weight should be 13-6 and not 9-1 (p. 430), but with the 
former atomic weight it would fall between caibon and nitrogen, and would 
be entirely out of place. Again, the element tellurium ought to come before 
iodine in the tables, but a recent determination of its atomic weight has con¬ 
firmed the number 128. And, finally, there ought not to be so many undis¬ 
covered or missing elements between didymium (145) erbium (166). 

The periodic law is, therefore, a valuable but still imperfect representation 
of certain chemical facts, and we must be careful not to accept the values for 
the atomic weights to which it points unless confirmed by evidence derived from 
other sources. It must for the present be still regarded as a hypothesis, until 
it has been more fully developed and until the anomalies which it contains have 
been explained. 


586 


Text-Book of Inorganic Chemistry. 


TABLES. 


I.— Weights and Measures. 

o 


i lb. Avoirdupois . 

i oz. Avoirdupois . 
i lb. Troy 

i oz. Troy 

x gallon.... 

i pint .... 
i gallon of water at 62° F 
( i 6°7 C.) weighs . 

1 cubic foot of water at 62° F 
weighs . 


16 oz. Av. 

7,000 grains 
437'5 grains 
12 oz. Troy 
5760 grains 
480 grains 
8 pints 

277*25 cubic inches 
20 fluid oz. 

70,000 grains 

62'326 lb. Av. 


1 metre (m ) . 


1 foot 
1 inch . 

1 litre (1.) 

1 pint 
1 fluid oz. 

1 kilogramme (kg.) 

1 gramme 


1 lb. Avoirdupois . 
1 oz. Avoirdupois . 
1 grain . 


= 10 decimetres (dm.) 

= 100 centimetres (cm.) 

= 1000 millimetres (mm.) 

= 39'37 inches 

= 30'48 centimetres 

- 25’40 millimetres 

= 1000 cubic centimetres (c.c.) 

= 176 pints 

= 568 cubic centimetres 

= 28’4 ,, 

= i.ooo grammes (g.) 

= 2’205 lb. Av. 

= 10 decigrammes (dg.) 

= 100 centigrammes (eg.) 

= 1000 milligrammes (mg.) 

= 15'43 2 grains 

— 453'6 grammes 

= 28*35 

= 64*8 milligrammes 




Appendix. 587 

II.— Weight of one cubic cetitimetre of Dry Air in milli¬ 
grammes at varying temperatures and pressures. 


(Temperature in ° C. and Pressure in mm. of Mercury.) 


t 

/=745 

750 

755 

760 

765 

770 

775 

0 

0 

1*268 

1*276 

1*285 

1*293 

1*301 

1*310 

I *3 I 9 

I 

1*264 

1*271 

1*280 

1*288 

1*296 

1*305 

i*3M 

2 

1*259 

1*267 

1*275 

1*283 

1*292 

1*300 

i * 3 i0 

3 

1*254 

1*262 

1*271 

1*279 

1*287 

1*296 

1*305 

4 

1*250 

1*258 

1*266 

i *275 

1*283 

I*29I 

1.300 

5 

1*245 

1*253 

1*261 

1*270 

1*2/8 

1*286 

1*295 

6 

1*240 

1*248 

1 *256 

1*265 

1*273 

I*28l 

1*290 

,7 

1*236 

1*244 

1*252 

1*261 

1*269 

1*277 

1*286 

8 

1*232 

1*240 

1*248 

1*256 

1*264 

1*273 

1*282 

9 

I *228 

1*236 

1*244 

1*252 

1*260 

I *268 

1*277 

10 

1*223 

1*231 

1*239 

1*247 

i’25S 

1*263 

1*272 

11 

1*219 

1*227 

1*235 

i*243 

1*251 

1*259 

1*268 

12 

1*215 

1*223 

1*231 

1*239 

1*247 

1*255 

1*263 

13 

I*2IO 

1*218 

1*226 

1*234 

1*242 

I*250 

1*259 

14 

1*206 

1*214 

1*222 

1*230 

1*238 

1*246 

i*254 

15 

1*202 

1*210 

1*218 

1*226 

1*234 

I*242 

1*250 

16 

1*198 

1*206 

1*214 

1*222 

1*230 

1*238 

1*246 

17 

1*194 

1*202 

1*210 

1*218 

1*226 

1*234 

1*242 

18 

1*189 

1*197 

1*205 

1*213 

1*221 

1*229 

1*237 

19 

I*l85 

1*193 

1*201 

1*209 

1*217 

1*225 

!*233 

20 

1*181 

1*189 

I*I 9 7 

1*205 

1*213 

1*221 

1*229 

21 

1*177 

1*185 

I*I 93 

1*201 

1*209 

1*217 

1*225 

22 

1 '173 

i*i8i 

1*189 

1*197 

1*205 

1*213 

1*221 

23 

1*169 

1*176 

1*184 

1*192 

1*200 

1*208 

1*216 

24 

1*165 

1*172 

1*180 

1*188 

1*196 

1*204 

I *212 

25 

1 *161 

1*168 

1*176 

1*184 

1*192 

1*200 

1*208 


HI._ Weight of one cubic centimetre and volume of one gramme 

of Water at varying temperatures. 

(Mean of Several Determinations.) 


Temp. 

Weight of 1 c.c. 
ih grammes 

Volume of i gr. 
in c.c. 

Temp. 

Weight of 1 c.c. 
in grammes 

Volume of i gr. 
in c.c. 

0 

0 

0*999878 

1*000122 

0 

13 

0*99943° 

I *000570 

1 

0*999933 

1*000067 

14 

0*999297 

I *000703 

2 

0*999972 

1 *000028 

15 

°* 999 I 54 

I *000847 

3 

0*999993 

1*000007 

16 

0*999004 

I *000997 

4 

1 *000000 

1*000000 

17 

0*998839 

1*001162 


0*999992 

1 *000008 

18 

0*998663 

1*001339 

j 

6 

0*999969 

1*000031 

19 

0*998475 

1*001527 

7 

0 ’999933 

1 *000067 

20 

0*998272 

1*00x731 

8 

0*999882 

i*ooon8 

21 

0*998065 

1 *001939 

Q 

0*999819 

1 *000181 

22 

0*997849 

I *002156 

y 

10 

0*999739 

1 *000261 

23 

0*997623 

1*002383 

11 

0*999650 

1 *000350 

24 

0*997386 

1*002621 

12 

o ’999544 

1 *000456 

25 

0*997140 

I *002868 








































5^8 


Text-Book of Inorganic Chemistry. 


IV .—Tension of Aqueous Vapours at varying temperatures. 


(Tension in mm. of Mercury.) 


Temp. 

Tension 

Temp. 

Tension 

Temp. 

Tension 

Temp. 

Tension 

o 

-15 

1'44 

0 

+ 7 

7'47 

O 

+ 17 

14*4° 

O 

+ 5° 

91*98 

—10 

2-15 

8 

7'99 

18 

I 5'33 

60 1 

148*88 

I “ 5 

3-16 

9 

8'55 

19 

16*32 

7° 

233'3l 

0 

4'57 

10 

9‘ r 4 

20 

17*36 

80 

354'87 

! + I 

4’9 I 

11 

977 

21 

18 *47 

90 

525'47 

| 2 

5’ 2 7 

12 

IO ‘43 

22 

19*63 

100 

760*00 

3 

5-66 

13 

11*14 

23 

20*86 

no 

1075'4 

4 

6'07 

14 

ii*88 

24 

22*15 

120 

i49i'3 

S 

6-51 

15 

122*67 

25 

2 3'5 2 

130 

2030*3 

6 

6-97 

16 

i3'Si 

26 

24*96 

140 

2717*6 


V.— Boiling-point of Water (/) under varying barometric pressure 

(£)■ 


A 


A 

t. 

A 

t. ■ 

A 

t. 


0 


0 


0 


0 

740 

99.26 

749 

99*59 

758 

99*93 

767 

100*26 

74i 

99*30 

750 

99-63 

759 

99*96 

768 

100*29 

742 

99'33 

75i 

99'6 7 

760 

100*00 

769 

100*33 

743 

99'37 

752 

9971 

761 

100*04 

770 

100*36 

744 

99*41 

753 

9974 

762 

100*07 

771 

100*40 

745 

99*44 

754 

99*78 

763 

100*11 

772 

100*44 

746 

99*48 

755 

99*82 

764 

100*15 

773 

100*47 

747 

99'52 

756 

99'85 

765 

100*18 

774 

100*51 

778 

99*56 | 

757 

99*89 

766 

100*22 

775 

100*55 


VI .—Corrections to reduce the Readmgs of a Baro 7 neter with a 
Glass Scale at different temperatures to o°. (Landolt and 

Bornstein .) 


(The Corrections are in mm., and are to be deducted from the Readings.) 


Temp. 

740 mm. 

750 mm. 

760 mm. 

770 mm. 

780 mm. 

0 

1 

0*13 

0*13 

0*13 

0*13 

0*13 

2 

0*26 

0*26 

0*26 

0*27 

0*27 

3 

0*38 

o*39 

o'39 

0*40 

0*40 

4 

0*51 

0*52 

o*53 

°*54 

°'54 

5 

0*64 

0*65 

o*66 

0*67 

0*67 

6 

0 77 

0*78 

079 

o*8o 

o*8i 

7 

0*90 

0*91 

0*92 

°’93 

o'94 

8 

1*02 

1*04 

1*05 

1 1*07 

1*08 






























































Appendix. 


589 


VI.— Continued. 


Temp. 

740 mm. 

750 mm. 

760 mm. 

770 mm. 

780 mm. 

O 

9 

I ’ I 5 

ri 7 

I*i8 

1*20 

1*21 

10 

1’28 

1'30 

I ’ 3 I 

i '33 

i '35 

11 

1 * 4 ! 

i ‘43 

!’45 

i '47 

1 * 4 8 

12 

!'54 

1*56 

158 

1'60 

1*62 

13 

1 *66 

1*69 

171 

173 

i '75 

14 

179 

1*82 

1*84 

1*87 

1*89 

IS 

I '02 

i *95 

1*97 

2*00 

2*02 

I 16 

2*05 

2*o8 

2*10 

2*13 

2*16 

1 x 7 

2 ’l8 

2*21 

2*24 

2*26 

2*29 

18 

2’30 

2 '34 

2 '37 

2*40 

2'43 

19 

2*43 

2*47 

2*50 

2'53 

2*56 

20 

2-56 

2*60 

2*63 

2*66 

2*70 

21 

2*69 

2*72 

2*76 

2*80 

2*83 

22 

2*82 

2*85 

2*89 

2 '93 

2*97 

23 

2-94 

2*98 

3*02 

3*06 

310 

i 2 4 

3'°7 

3 'n 

3*16 

3 '20 

3'24 

! 2 5 

3 '20 

3 '24 

3' 2 9 

3'33 

3'37 


VII .—Specific Gravity and Composition of Dilute and 
Concentrated Acids. 


Sulphuric Acid 

Nitric Acid 

Acetic Acid 

Hydrochloric Acid 

Per cent. 

H 2 SO* 

by weight j 

Specific 

gravity 

Per cent. 1 
HNO, 1 
by weight 

Specific 

gravity 

Per cent. 

hc 2 h ,o 2 

by weight 

Specific 

gravity 

Per cent. 
HC1 

by weight 

Specific 

gravity 

3 

1*031 

5 

I *030 

5 

1*0067 

5 

I *027 

10 

1*069 

10 

1*059 

10 

1*0142 

10 

1*050 

15 

i*io6 

15 

1*089 

iS 

I *0214 

15 

1-075 

20 

1 *146 

20 

1*120 

20 

I *0284 

20 

1*101 

25 

1*184 

25 

i'i 53 

25 

1 '0350 

25 

1*126 

3 ° 

1*224 

30 

1*185 

30 

1 *0412 

3 ° 

I-I 5 I 

35 

1*266 

35 

1 *218 

35 

1 ’0470 

35 

1*177 

40 

1*306 

40 

1*251 

40 

1*0523 

40 

1*200 

45 

1*051 

45 

1*284 

45 

1-0571 

42*4 

1*210 

5 ° 

i *399 

50 

i* 3 T 9 

50 

1*0615 



55 

1*448 

55 . 

1*346 

55 

l '0653 



60 

1*502 

60 

i '374 

60 

1 *0685 



65 

i'SS 7 

65 

i *395 

65 

1*0712 



70 

1*615 

70 

1*423 

70 

i '0733 



75 

1 ‘675 

75 

1-442 

75 

1 *0746 



80 

1-733 

80 

1*460 

80 

1 *0748 



85 

i' 7 8 3 

85 

1-478 

85 

1*0739 



9 ° 

1*819 

90 

i '495 

90 

1*0713 



95 

1*839 

95 

i‘ 5 T 3 

95 

1 *0660 



100 

1*838 

100 

i' 53 ° 

100 

1 i'o 553 

—•—* — 

■ 



















































590 Text-Book of Inorganic Chemistry. 

VIII.— Specific Gravity a7id Compositicni of Alkaline Solutions. 


Caustic Potash 

Caustic Soda 

Ammonia 

Per cent. 
KOH 

Specific 

Per cent. 
NaOH 

Specific 

Per cent. 
NH, 

Specific 

by weight 

gravity 

by weight 

gravity 

by weight 

gravity 

s 

I '041 

5 

1-058 

5 

0-979 

10 

1-083 

10 

I-115 

10 

°’959 

15 

1-128 

15 

1-170 

15 

0-941 

20 

1-177 

20 

1-225 

20 

0-925 

25 

1-230 

25 

1-279 

25 

0911 

3 ° 

1*288 

30 

i '332 

3 ° 

0-898 

35 

i '349 

35 

1'384 

33 

0-891 

40 

1-412 

40 

1 '437 

3 6 

0884 

45 

1 ‘475 

45 

1-488 



50 

1 ’539 

50 

1 '540 



55 

1-604 

55 

1 ' 59 i 



60 

1-667 

60 

1'643 



65 

1729 

65 

1 '695 



70 

1-790 

70 

1748 




To convert a density expressed in Twaddclis degrees into true specific 
gravity :— 

Divide by 200 and add 1. Thus, 25 0 Twaddell = + i = i , i25. 


IX .—Heat of Formation a?id Solution of some Important 
Compounds. (After Thomsen.) 

Note. —The quantity of heat (H) is expressed in the number of kilogrammes 
of water which would be raised i° C. in temperature on the formation of the 
molecular weight of the compound in grammes. Thus : — 

(H 2 , O) = + 68-4 

means that on the formation of 18 grammes of water, or on the combustion of 2 
grammes of hydrogen, the quantity of heat set free would raise 68-4 kilos, of 
water i°. If heat is absorbed, instead of liberated, the number is expressed as 
a negative quantity. 


Compound 

H 

Remarks 

(H-\ 0 ) . . 

+ 68-4 

Liquid water 

(H 2 , 0 2 , aq.) . 

+ 45'3 

Formation of H.,Oo in dilute solution 

(H 2 ,S) . 

+ 47 

Gaseous HoS 

(ITS, aq.) 

+ 46 

Solution of H.>S in water 

(H.Cl) . 

+ 22 ’O 

Gaseous HC 1 

(HC 1 , aq.) 

+ 17-3 

Solution of HC 1 in water 

(H,Br) 

+ 8-4 

Gaseous HBr 

(HBr, aq.) 

+ iQ "9 

Solution of HBr in water 

(H,I) . . 

— 6"o 

Gaseous HI 

(HI, aq.) . 

+ 19-2 

Solution of HI in water 

(H.Cl.O 5 , aq.) . 

+ 23-9 

Formation of HC 10 3 in dilute solution 

(H,Br, 0 3 , aq.) . 

+ 12-4 

HBr 0 3 
























Appendix. 591 


IX.— Continued. 


Compound 

H 

Remarks 

(H, I.O 3 , aq.) . 
(S.O 2 ) 

(SO 2 , aq.). 

(S. 05 ) . . 

(SO, aq.). 
(H 2 ,S, 0 4 ) . 
(H 2 S 0 4 , aq.) . 
(N,H 5 ) . 

(NH 5 , aq.) 
(N,H 4 ,C 1 ) 

(N 2 . 0 ) . 

(N.O) • 

(H,N, 0 2 , aq.) . 

(N.O 2 ) . 

(H.N.O 3 ) . 
(HNO'\ aq.) . 
(F'.O 6 ) • 

(H 5 ,P, 0 4 ). 

(P,C 1 5 ) . 

(P.C 15 ) . 

(As, Cl 5 ) . 
(Sb.Cl 5 ) . 

(C,0) . . 

(C.o 2 ) . . 

(C,H 4 ) . . 

(C 2 , H 4 ) . 

(C 2 , H 2 ) . 

(C 2 N 2 ) 

(K,C 1 aq.) 
(K,Br, aq.) 

(K,I, aq.) 
(K,C 1 , 0 5 ) . 
(K 4 S 0 4 ) . 
(K 2 S 0 4 ,H-S 0 4 ). 
(K.N.O 5 ) . 
(KN0 3 , aq.) . 
(Na,Ci, aq.) 
(Ca,Cl 2 , aq.) . 
(Ca.C.O 5 ). 
(Ba.Cl-, aq.) 
(Ba.SO 4 ) . 
(Al-'Cl 8 ) . 

(APC 1 6 , aq.) . 

(Fe 2 ,Cl 6 . aq ) . 

(Fe,Cl 2 , aq.) . 
(Zn,S, 0 4 , aq.) . 
(Cu.O) . 
(Cu.S.O 4 ). 
(Cu.S.O 4 , aq.) . 

(Ag.Cl) . . 

(Ag.Br) . • 

(Ag.I) • • 

(Ag N.O 3 , aq.) . 

+ 55 ' 8 

+ 7i’i 

+ 77 

+103'2 
+ 39-2 
+192 '9 
+ i 7'9 
+ ji 'g 
+ 8 '4 

+ 75 '8 

- I 7‘5 

- 21 ‘6 

+ 3 °' 8 

- 2 ‘0 
+ 41"6 

+ 7*5 

+ 369-9 
+ 302 '6 
+ 105-0 
+ 75‘3 
+ 7 1 4 
+ 9 1 '4 
+ 29-0 
+ 97-0 
+ 21'8 

- 27 

- 48-2 

- 657 

+ IOI "2 
+ 90*2 
+ 75 
+ 95'9 
+ 344 6 

+ i6"6 

+ H 9‘5 

- 8‘5 
+ 9 6 '5 
+ 187-2 
+ 270*4 
+ 196-8 

+ S 3 8 ’ 1 
+ 322 0 

+ T 53’7 
+ 255 ‘4 

+ IOO’O 

+ 248 
+ 37-2 
+ 182-6 
+ 198-4 
+ 29-4 
+ 22*7 
+ I 3 ' 8 
+ 237 

Formation of HI0 5 in dilute solution 

Gaseous SOo 

Solution of SOo in water 

Liquid SO 3 

Solution of S0 3 in much water 

Liquid H 0 SO 4 

Solution of H 0 SO 4 in much water 

Gaseous NH 5 

Solution of N H 3 in water 

Solid NH 4 CI 

Gaseous NoO 

Gaseous NO 

Formation of HN0 2 in dilute solution 

Gaseous N0 2 

Liquid HNO* 

Solution of HNO 3 in water 

Solid P 0 O 5 

Crystalline H 3 P0 4 

Solid PCI 5 

Liquid PCI 3 

Liquid AsC1 3 

Liquid SbCl 3 

Gaseous CO from amorphous C 
,, COo 

Gaseous CH 4 
,, C 0 H 4 

„ C 2 Ho 

(CN), 

Formation of KC1 in dilute solution' 

„ KBr 

K1 

Crystalline KCIO 3 

,, k ? so 4 

Formation of KHS0 4 from K 0 SO 4 and H 2 S0 4 
Crystalline KN0 3 

Solution of KNO 3 in water 

Formation of NaCl in dilute solution 
,, CaCl 2 

CaC0 3 as calcite 

Formation of BaCL in dilute solution 

BaS0 4 as an amorphous powder 

Anhydrous A1 2 C1 6 (solid) 

Solution of A1 2 C1 6 in water 

Formation of Fe ? Cl 6 in dilute solution 

FeCl 2 
,, ZnS0 4 

Solid CuO 

Anhydrous CuS0 4 (solid) 

Formation of CuS0 4 in dilute solution 

Solid AgCl 
„ AgBr 
,, Agl 

Formation of AgN0 3 in dilute solution 
























INDEX. 


ACI 

Acids. 56 
Acid hydrates, 56 
Affinity, 45, 51 
Agate, 269 
Air, 200 
Air bath, 89 
Alabandite, 463 
Alabaster, 394 
Albite, 365, 425 
Albumen, 140 
Algaroth powder, 257 
Alkali manufacture, 370 

— waste, 169, 372 
Alloys, 58 
Alum, 423 

— cake, 423 
- shale, 424 

Alumina, 421 
Aluminates, 422 
Aluminic acid, 422 
Aluminite, 423 
Aluminium, 419 

— bronze, 531 

— chloride, 428 

— detection of, 429 

— fluoride, 428 

— gold, 531 

— hydrate, 421 

— oxide, 421 

— phosphate, 425 

— salts of, 422 

— silicate, 425 

— sulphate, 422 
. - sulphide, 428 

Alunite, 140, 424 
Amalgamation process, 543 


ANT 

Amalgams, 58, 541 
Amethyst, 269 
Amidogen, 183, 188 
Ammonia, 183 

— soda process, 372 
Ammonium, 183, t88 , 382 

— bromide, 386 

— carbonate, 383 

— chloride, 385 

— chlorplatinate, 562 

— chromate, 477 

— detection of, 387 

— dichromate, 477 

— ferrocyanide, 456 

— nitrate, 383 

— phosphates, 383 

— sulphate, 382 

— sulphide, 386 

— sulphydrate, 386 

— sulphostannate, 520 
Amorphous carbon, 282 
Analysis, 4 

Anatase, 322 
Analcime, 425 
Andalusite, 425 
Anglesite, 498, 503 
Anhydrides, 56 
Anhydrite, 394 
Animal charcoal, 283 
Anorthite, 425 
Anthracite, 284 
Antichlor, 103, 369 
Antimonic acid, 254 
— anhydride, 254 
—. chloride, 257 
Antimonite, 250, 258 

Q Q 



594 


Index. 


ANT 

Antimoniuretted hydiogen, 252. 
Antimonous chloride, 256 

— sulphate, 254 

— sulphide, 258 
Antimony, 250 

— compounds of, with oxygen, 25 3 
-the halogens, 256 

-sulphur, 258 

— glass, 259 

— pentachloride, 257 

— pentasulphide, 259 

— pentoxide, 254 

— sulphochloride, 257 

— tetroxide, 256 

— tribromide, 257 

— trichloride, 256 

— triiodide, 258 

— trioxide, 253 

— trisulphide, 258 
Apatite, 206, 396 
Appendix, 573 
Aquamarine, 429 
Aqua regia, 198 
Argentan, 492 
Argillaceous iron ore, 436 
Arragonite, 286, 399 
Arsenic, 229 

— anhydride and acid, 244 
Arsenical pyrites, 229, 437 
Arsenic bloom, 233 

— compounds of, with oxygen, 233 

— •— the halogens, 246 
-sulphur, 248 

— detection of, 237 

— disulphide, 248 

— pentasulphide, 250 

— trisulphide, 248 
Arsenious acid, 235 

— anhydride, 233 

— bromide, 247 

— chloride, 247 

— iodide, 247 

— fluoride, 248 

— sulphide, 248 
Arseniuretted hydrogen, 231 
Arsine, 231 

Asbestos, 413, 417 

— platinized, 162 
Atacamite, 528 
Atmospheric air, 200 


BIS 

Atomic heat, 579 

— weights, table of, 60 

— theory, 44 

— weights, determination of, 57 

-relations between, 65, 582 

Atomicity, 47 

Atoms, 45 
Augite, 413, 417 
Auric acid, 554 

— chloride, 553 

— cyanide, 554 

— oxide and hydrate, 554 

— sulphide, 555 
Aurous chloride, 555 

— cyanide, 555 

— oxide, 555 
Avogadro’s law, 50, 1575 
Azurite, 523, 528 


Barium, 408 
-— carbonate, 411 

— chloride, 411 

— chromate, 477 

— detection of, 412 

— hydrate, 409 

— nitrate, 411 

— oxide, 408 

— peroxide, 410 

— sulphate 411- 

— sulphide, 412 
Baryta, 408 
Baryta-water, 409 
Basic hydrates, 56 
Bay salt, 377 
Bell-metal, 531 
Beryl, 429 
Beryllium, 429 

— - carbonate, 430 

— chloride, 430 

— hydrate, 430 

— oxide, 430 

— sulphate, 430 
Bessemer process, 441 
Biotite, 413 
Bismuth, 513 

— chloride, 515 

— chromate, 515 

— detection of, 515 

— nitrate, 515 


Index . 


595 


BIS 

Bismuth oxide and hydrate, 514 

— suboxide, 514 

— sulphide, 515 
Bismuthic acid, 514 
Bismuthinite, 515 
Black ash, 371 

— lead, 281 
Blast furnace, 437 

Bleaching powder, 102, 116, 400' 

Blende, 140, 493, 496 

Blue vitriol, 525 

Bog iron ore, 436 

Bole, 427 

Bone ash, 396 

— black, 283 
Boracite, 260, 417 
Borax, 376 

Boric acid and anhydride, 262 
Boron, 260 

— nitride, 266 

— sulphide, 266 

— tribromide, 265 

— trichloride, 264 

— trifluoride, 265 
Boronatro-calcite, 260 
Boulangerite, 507 
Brass, 530 
Braunite, 460, 463 
Breithauptite, 250 
Britannia metal, 252, 521 
Bromic acid, 122 
Bromide of lime, 123 
Bromides, 122 
Bromine, 118 

Bronze, 531 
Brookite, 322 
Brown coal, 285 
Brucite, 415 
Burnt alum, 423 
Bunsen burner, 309 

— pump, 33 


Cadmium, 497 

— compounds of, 497 

— detection of, 498 
Caesium, 380 

— compounds of, 381 
Calamine, 493, 495 
Calcareous tufa, 400 


CHA 

Calcite, 286, 399 
Calcium, 388 

— carbonate, 398 

— chloride, 404 

— chromate, 477 

— detection of, 406 

— diacid phosphate, 398 

— fluoride, 405 

— hydrate, 390 

— monacid phosphate, 398 

— nitrate, 396 

— oxalate, 404 

— oxide, 389 

— pentasulphide, 405 

— peroxide, 393 

— phosphate, 396 

— silicate, 401 

— sulphate, 394 

— sulphide, 405 

— sulphydrate, 405 
Calomel, 540 
Calcspar, sec Calcite 
Cannel coal, 285 

Caput mortuum, 163, 447 
Carbamic acid, 293 
Carbon, 277 

— disulphide, 300 

— compounds of, 286 
-with hydrogen, 303 

— hexachloride, 310 

— oxychloride, 297 

— oxysulphide, 302 

— tetrabromide, 310 

— tetrachloride, 309 
Carbonic acid or anhydride, 286 

— oxide, 294 
Carbonyl chloride, 297 
Carnallite, 418 
Cassel yellow, 506 
Cassiterite, 516, 519 
Cast-iron, 438 
Caustic potash, 346 

— soda, 367 
Celestine, 406 
Cement copper, 526 

— steel, 441 
Cerite, 433 
Cerium, 433 
Cerussite, 498, 504 
Chalcedony, 269 


Q Q 2 


596 


Index. 


CUP 


CHA 

Chalcostibite, 250 
Chalk, 286, 400 
Chalybeate springs, 446 
Chameleon mineral, 466 
Charcoal, 282 
Chemical affinity, 4, 51 

— calculations, 61 

— combination, laws of, 37 

— elements, remarks on, 63 

— .equivalents, 42 

— nomenclature, 54 

— processes, 3 

— properties, 77 

—- symbols and formulae, 5$ 

Chili saltpetre, 369 

China-clay, 426 

Chlorchromic acid, 477 

Chloric acid, no 

Chloride of lime, 102, 116, 400 

Chlorides, 109 

Chlorine, 97 

— hydrate, 103 

— oxides and oxy-acids, 109 
peroxide, 114 

Chlorous anhydride and acid, 114 
Chlorplatinic acid, 561 
Chlorsulphonic acid, 165 
Chromates, 474 
Chrome alum, 425, 471 

— iron ore, 468 

— — red, 506 

— yellow, 506 

Chromic anhydride and acid, 473 

— chloride, 472 

— hydrate, 470 

— nitrate, 472 

— oxide, 469 

— phosphate, 472 

— sulphate, 471 
Chromites, 470 
Chromium, 468 

— detection of, 479 
Chromous chloride, 473 

— hydrate, 473 
Chromyl chloride, 478 
Chrysoberyl, 429 
Cinnabar, 533, 538 
Clay, 426 

Clay ironstone, 436, 446 
Coal, 284 


Coal gas, 319 
Cobalt and nickel, 483 
Cobalt, 484 

— arsenides, 489 

— bloom, 230 

— detection of, 489 
Cobaltic chloride, 486 

-ammoniacal compounds of, 

487 

oxide and hydrate, 485 
Cobaltous carbonate, 486 

— chloride, 487 

— cyanide, 488 

— nitrate, 486 

— nitrite, 486 

— oxalate, 487 

—- oxide and hydrate, 485 
phosphate, 486 

— sulphate, 486 

— sulphide, 488 
Coke, 284 
Colcothar, 163, 447 
Cold-short, 440 
Columbite, 330 
Common salt, 365, 376 
Combustion, 29 
Compound radical, 61 
Condy’s fluid, 467 
Contact action, 162 
Copper, 523 

— alloys of, 530 

— detection of, 531 
Copper-glance, 523, 530 
Copper-pyrites, 140, 437, 523, 530 
Coprolites, 206, 396 

Corrosive sublimate, 536 
Corundum, 421 
Couple, galvanic, 318 
Covellite, 529 
Crocoisite, 506 
Crocus antimonii, 259 
Cryolite, 134, 365, 420, 428 
Crystals, how obtained, 75 
Crystalline systems, 71 
Crystallography, 70 
Cubic alum, 423 

— nitre, 369 
Cupellation, 499, 551 
Cupric acetates, 528 

— ammonium sulphate, 526 




Index . 


597 


CUP 

Cupric arsenite, 527 

— bromide, 528 

— carbonates, 527 

— chloride, 5 2 & 

— compounds, 5 2 4 

— hydrate, 525 

— nitrate, 527 

— oxide, 524 

— phosphate, 527 

— sulphate, 525 

— sulphide, 528 
Cuprite, 523, 529 
Cuprous chloride, 5 2 9 

— compounds, 529 

— hydrate, 529 

— iodide, 530 
-oxide, 529 

— sulphide, 530 
Cyamelide, 317 
Cyanic acid, 316 
Cyanogen, 311 

— bromide, 319 

— chlorides, 318 

— iodide, 319 
Cyanuric acid, 317 

Dalton’s atomic theory, 45 

Dark-red silver ore, 250, 543, 551 

Datolite, 260 

Decrepitation, 87 

Deliquescence, 85 

Density of gases, 51 

Desiccator, 87 

Dialysed iron, 448 

Diamond, 279 

Diaspore, 422 

Dibasic acids, 56 

Dichromic acid, 47 5 

Didymium, 433 

Dimorphism, 75 

Displaceable hydrogen, 56 

Diselenium dichloride, 179 

Disthene, 425 

Disulphur dichloride, 1 73 

Disulphuric acid, 166 

Dithionic acid, 170 

Divalent elements, 47 

Dolomite, 286, 413, 4 X 7 

Drummond’s light, 25 


FLA 

Dulong and Petit’s law, 70, 578 
Dutch metal, 531 
Dyad, 47 
Dyscrasite, 250 


Earthenware, 426 
Eaux de Javelle, 118, 355 
Electro-chemical theory, 46 
Electrodes, 6 
Electro-gilding, 556 
Electrolysis, 6 
Electro-plating 550 
Elements, 5 

— general remarks on, 63 
Emerald, 429 

Emery, 421 
Epsom salt, 140, 415 
Equivalent weights, 42 
Erbium, 433 
Etching on glass, 136 
Ethylene, 306 

— dibromide, 307 

— dichloride, 307 
Eudiometer, 25 


Fahlerz, 523 
Fayence, 426 
Felspar, 420, 425 
Ferric acid, 450 

— chloride, 451 

— disulphide, 452 
—hydrate, 447 

— nitrate, 449 

— oxide, 447 

— phosphate, 449 

— sulphate, 448 
Ferromanganese, 439 

Ferrous and ferric compounds, 444 

— carbonate, 446 

— chloride, 450 

— hydrate, 445 

— iodide, 451 

— oxide, 445 

— phosphate, 447 

— sulphate, 445 

— sulphide, 452 
Fining, 439 

Flame, structure of, 30b 


598 


Inaex. 


IRO 


FLI 

Flint, 269 

Flowers of sulphur, 142 
Fluoboric acid, 265 
Fluorine, 133 
Fluor-spar, 134, 388, 405 
Fluosilicic acid, 275 
Fluotitanic acid, 325 
Fuller’s earth, 427 
Fulminating gold, 554 
— silver, 546 

Fuming sulphuric acid, 163 
Fur of kettles, 83 
Fusible metal, 522 


Gadolinite, 433 
Galena, 140, 498, 507 
Gahnite, 422 
Gallium, 431 
Galvanic couple, 518 
Galvanized iron, 443, 494 
Gasometer, 15 
Gay-Lussac’s law, 43 
German silver, 492, 531 
Gibbsite, 421 
Glance cobalt, 484, 489 
Glass, 401 

Glauber’s salt, 140, 367 
Glucinum, sec Beryllium 
Gold, 552 

— alloys of, 556 

— compounds of, 553 
Grain-tin, 517 
Granite, 269, 420 
Graphic tellurium, 552 
Graphite, 281 
Graphitic acid, 282 
Green vitriol, 445 
Greenockite, 497 
Guignet’s green, 471 
Gun-metal, 531 
Gunpowder, 351 
Gypsum, 140, 388, 394 


Hematite, 436, 447 
Halogens, 96 
Haloid-acids, 57 
Haloid salts, 57, 96 
Hard water, 83 


Hartshorn, spirits of, 183 
Hauerite, 463 
Hausinannite, 460, 464 
Heavy carburetted hydrogen, 304. 
Heavy metals, 436 
Heavy spar, 140, 408 
Hexads, 47 
Hornblende, 413, 417 
Horn-silver, 543 
Hyacinth, 434 
Hydrargillite, 421 
Hydrates, 56, 87 
Hydraulic mortar, 393 
Hydriodic acid, 127 
Hydrobromic acid, 120 
Hydrochloric acid, 104 
Hydrocobalticyanic acid, 488 
Hydrocyanic acid, 313 
Hydroferricyanic acid, 458 
Hydroferrocyanic acid, 457 
Hydrofluoric acid, 134 
Hydrogen, 16 
Hydrogen peroxide, 90 
— persulphide, 148 
Hydrosulphurous acid, 155 
Hydroxylamine, 186 
Hypobromous acid, 123 
Hypochlorous anhydride and acid 
116 

Hypophosphoric acid, 220 
Hypophosphorous acid, 222 
Hyposulphuric acid, 170 
Hyposulphurous acid, 167 


Ice, 80, 81 
Iceland-spar, 399 
Imidogen, 183, 188 
Indium, 432 

— compounds of, 432 
Infusorial earth, 269 
Iodic acid, 129 

— anhydride, 130 
Iodides, 129 
Iodine, 123 

— chlorides, 133 

— oxy-acids, 129 
Iridium, 568 

chloride, 569 
Iron, 436 


Index. 


599 


IRO 

Iron alum, 425, 448 

— compounds, 444 

— compounds of, with cyanogen, 

453 . 

— detection of, 459 

— pyrites, 140, 437, 452 

— rust, 442 

—- spinelle, 422 
Isomeric compounds, 38 
Isomorphism, 75 


Jamesonite, 507 
Jasper, 269 
Jet, 285 


Kaolin, 426 
Kermes mineral, 259 
Kieselguhr, 269 
Kieserite, 415 
Kupferniclcel, 229, 490 


Lakes, 422 
Lampblack, 285 
Lanthanum, 433 
Lapis lazuli, 427 
Lavoisier, 30, 31 
Law, Avogadro’s, 50, 575 
c 1 — Gay-Lussac’s, 43 

— Dulong and Petit’s, 70, 578 

— of isomorphism, 75, 580 
Lead, 498 

— acetate, 505 

— alloys of, 507 

— carbonates, 504 

— chloride, 506 

— chromate, 5°6 

— cyanide, 507 

— detection of, 508 

— hydrate, 5 02 

— iodide, 506 

— nitrate, 504 
—- oxide, 5 01 

— peroxide, 502 

— phosphate, 504 

— salts of, 503 

— suboxide, 501 

— silicates, 504 


MAG 

Lead sulphate, 503 

— sulphide, 507 
Lepidolite, 379 

Light, influence of, on chemical 
affinity, 53 

— carburetted hydrogen, 304 

— red silver ore, 543, 551 
Lignite, 285 

Lime, 389 

— chloride of, 102, 116, 400 

— felspar, 425 

— light, 25 

— milk of, 390 

— water, 390 
Limestone, 286, 399 
Limonite, 436 
Linnaeite, 489 
Liquor ammonite, 183 
Litharge, 501 , 

Lithium, 378 

— carbonate, 380 

— chloride, 380 

— detection of, 380 

— oxide and hydrate, 379 

— phosphate, 379 

— sulphate, 379 
Lithographic stone, 399 
Liver of sulphur, 363 
Lunar caustic, 547 
Luteocoballic chloride, 487 


Magnesia, 414 

— alba, 417 

— calcined, 415 

— usta, 415 

Magnesite, 286, 413, 417 
Magnesium, 413 

— arsenate, 416 

— borate, 417 

— carbonate, 417 

— chloride, 417 

— detection of, 418 

— hydrate, 415 

— oxide, 414 

— phosphates, 416 

— silicate, 417 

— sulphate, 415 

Magnetic oxide of iron, 436, 449 

— pyrites, ^37, 453 


6 oo 


Index. 


MAL 

Malachite, 523, 528 
Manganese, 460 

— alum, 425, 463 

— detection of, 467 

— disulphide, 463 

— oxides, 461 

— perhydrate, 465 

— peroxide, 464 

— spar, 460, 462 
Manganic acid, 465 

— chloride, 464 

— hydrate, 463 

— oxide, 463 

— sulphate, 463 
Manganite, 460, 463 
Manganous carbonate, 462 
-— chloride, 462 

— nitrate, 462 

— oxide and hydrate, 461 

— sulphate, 461 

— sulphide, 463 
Marble, 286, 399 
Marcasite, 453 
Marking-ink, 547 
Marl, 427 
Marsh gas, 304 
Marsh’s test, 239 

Mass, influence of, on chemical 
affinity, 53 
Massicot, 501 
Matches, lucifer, 210 
Matlockite, 506 
Meerschaum, 413, 4T7 
Mercuric chloride, 536 

— chromate, 536 

— compounds, 535 

— cyanide, 538 

— iodide, 538 

— nitrate, 536 

— oxide, 535 

— sulphate, 535 

— sulphide, 538 
Mercurous chloride, 540 

— chromate, 540 

— compounds, 539 

— iodide, 541 

— nitrate, 539 

— oxide, 539 

— sulphide, 541 
Mercury, 533 


NIC 

Mercury, detection of, 542 
Metaboric acid, 263 
Metals, 331 

— classification of, 340 

— of the alkalies, 342 

— of the alkaline earths, 388 

— of the earths, 419 

— heavy, 436 

— noble, 533 

— platinum, 557 
Metantimonic acid, 255 
Metaphosphoric acid, 220 
Metastannic acid, 519 
Methane, 304 

Mica, 420 

Microcosmic salt, 383 
Milk of lime, 390 

— sulphur, [45 
Millerite, 492 
Mimetesite, 230 
Mineral waters, 82 
Minium, 502 
Mispickel, 229 

Molecular weights, how found, 51, 

573 

Molecule, 48 
Molybdenite, 325 
Molybdenum, 325 

— chlorides, 327 

— oxides, 326 

Molybdic anhydride and acid, 326 
Monads, 47 
Monobasic acids, 56 
Monovalent elements, 47 
Mortar, 391 

— hydraulic, 393 
Mosaic gold, 521 
Mottramite, 329 

Multiple proportions, law of, 42 
Muntz-metal, 531 


Nascent state, 50 

Native, 63 

Natrolite, 425 

Nessler’s solution, 387, 538 

Niccolite, 490 

Nickel, 490 

— alloys of, 492 

— carbonate, 491 


Index. 


601 


NIC 

Nickel chloride, 491 

— cyanide, 491 

— detection of, 492 

— nitrate, 491 

— sulphate, 491 

— - sulphide, 492 

— silver, 492 

Nickelic oxide and hydrate, 491 

Nickelous oxide and hydrate, 496 

Niobium, 330 

Nitrated acid, 158 

Nitrates, 192 

Nitre, 350 

Nitric acid, 189 

— anhydride, 192 

— oxide, 192 

— peroxide, 196 
Nitrites, 195 
Nitrogen, 181 

— compounds of, with oxygen, 189 

— chloride, 204 

— iodide, 206 

Nitrosulphonic acid, 166, 197 
Nitrosyl chloride, 198 
Nitrous anhydride and acid, 195 

— oxide, 198 
Nitryl chloride, 197 
Noble metals, 533 
Nomenclature, chemical, 54 
Nordhausen sulphuric acid, 163 


Olefiant gas, 306 
Olivine, 271, 413, 417 
Opal, 269 
Orangite, 434 
Orpiment, 229, 248 
Orthoclase, 272, 425 
Orthophosphoric acid, 216 
Osmic acid, 570 
Osmiridium, 558, 569 
Osmium, 570 
Oxalic acid, 297 
Oxamic acid, 299 
Oxamide, 300 
Oxidation, 36 
Oxides, 55 
Oxy-acids, 56 
Oxyammonia, 186 
Oxyammonium chloride, 387 


PHO 

Oxyammonium sulphate, 387 
— salts, 386 
Oxychlorides, 58 
Oxygen, 8 

Oxy-hydrogen blowpipe, 24. 
Ozone, 91 


Palladic chloride, 568 

— oxide, 568 
Palladium, 566 

— hydride, 566 
Palladous.chloride, 56S 

— iodide, 568 

— nitrate, 567 

— oxide, 568 
Paracyanogen, 312 
Parkes’s process, 5C0 
Passive iron, 443 
Pattinson’s process, 500 
Pearlash, 353 

Peat, 285 
Pentads, 47 

Pentavalent elements, 47 
Pentathionic acid, 172 
Perchloric acid, 113 
Perchromic acid, 474 
Periodic acid, T31 

— law, 583 

Permanganic anhydride and acid, 
466 

Peroxides, 55 
Persnlphocyanic acid, 318 
Petallite, 379 
Phenacolite, 230 
Phenakite, 429 
Phlogiston theory, 31 
Phosgene gas, 297 
Phosphates, 217 
Phosphine, 211 

Phosphonium iodide, 128, 213 
Phosphoretted hydrogen, 211 
Phosphoric acid, 216 

— anhydride, 214 

— bromide, 227 
—- chloride, 225 

— oxybromide, 227 

— oxychloride, 226 

— sulphide, 229 
Phosphorite, 206, 396 


602 


Index. 


PHO 

Phosphorous anhydride and acid, 
221 

— bromide, 227 

— chloride, 224 
Phosphomolybdic acid, 327 
Phosphotungstic acid, 328 
Phosphorus, 206 

— compounds of, with hydrogen, 

211 

-oxygen, 213 

-— the halogens, 223 

-- sulphur, 228 

— diiodide, 228 

— disulphide, 229 

— pentabromide, 227 

— pentachloride, 225 

— pentasulphide, 229 

— tribromide, 227 

— trichloride, 224 

— triiodide, 228 
Phosphoryl chloride, 226 
Photography, 549 
Physical processes, 3 

— properties, 67 
Pink salt, 520 
Pitchblende, 479 
Plaster of Paris, 394 
Platinic chloride, 561 

— oxide and hydrate, 562 

— sulphide, 562 
Platinous chloride, 562 

— cyanide, 563 

— oxide and hydrate, 563 

— sulphide, 563 
Platinum, 558 

— ammoniacal compounds of, 563 

— black, 559 

— metal', 557 

— spongy, 558 
Plumbago, 281 
Poles or electrodes, 6 
Polythionic acids, 171 
Porcelain, 426 
Portland cement, 393 
Potash, 353 

— alum, 423 

— felspar, 425 
Potassium, 342 

— acid carbonate, 354 

— acid fluoride, 358 


PRO 

Potassium acid sulphate, 349 

— aluminate, 422 

— aurate, 554 
—- bromide, 356 

— carbonate, 353 

— chlorate, 354 

— chlorchromate, 477 

— chloride, 356 

—- chlorpalladinate, 568 

— chlorplatinate, 364, 561 

— chromate, 476 

— cobalticyanide, 488 

— compounds of, 346 

— cyanate, 361 

— cyanide, 359 

— detection of, 364 

— dichromate, 475 

— ferrate, 450 

—- ferricyanide, 457 

— ferrocyanide, 453 

— fluoborate, 358 
fluoride, 358 

— fluosilicate, 359 

— hydrate, 346 

— hypochlorite, 354 

— iodide, 356 

— iridium chloride, 559 

— manganate, 465 

— monosulphide, 362 

— nitrate, 350 

— nitrite, 352 

— osmate, 571 

— oxalate, 355 

— oxide, 346 

— perchlorate, 354 

— permanganate, 466 

— peroxide, 346 

— plumbate, 507 

— polysulphides, 363 

— silicate, 355 

— stannate, 520 

— sulphate, 349 

— sulphocyanate, 361 

— sulphydrate, 363 

— uranate, 482 
Prehnite, 272 
Priestley, 8 

Proportions, combining, 38 

Proustite, 543 

Prout’s hypothesis, 582 


PRU 


Index . 


603 


Prussian blue, 456 
Prussic acid, 313 
Psilomelane, 460 
Puddling, 439 
Purple copper ore, 523, 530 

— of Cassius, 553, 556 
Pyrargyrite, 250, 543 
Pyrites, copper, 140, 523, 530 

— iron, 140, 437, 452 
Pyroboric acid, 263 
Pyrolusite, 460, 464 
Pyromorphite, 504 
Pyrophosphoric acid, 219 


Quantity (mass), influence of, on 
chemical affinity, 53 
Quartz, 268 
Quicklime, 389 
Quicksilver, see Mercury 


Radical, compound, 61 
Rain water, S3 
Realgar, 229, 248 
Red arsenic glass, 248 

— copper ore, 523, 529 

— lead, 502 

— prussiate, 457 

— short, 440 
Reduced iron, 442 
Reduction, 36 
Rhodium, 570 
Rock salt, 365, 376 
Roman alum, 424 

— cement, 393 
Roseocobaltic chloride, 487 
Rouge, 447 

Rubidium, 380 

— compounds, 381 
Ruby, 421 

— sulphur, 248 
Ruthenium, 571 
Rutile, 322 


Sal-ammoniac, 385 
Salt-cake, 371 
Saltpetre, 350 
Salts, constitution of, 332 


SIL 

Salts of sorrel, 355 
Samarskite, 330, 433 
Sandstone, 269 
Sassolite, 260 
Scheele, 8, 31, 97 
Scheele’s green, 235, 527 
Scheelite, 327 
Schlippe’s salt, 259 
Schiitzenberger’s acid, 155 
Schweinfurth green, 235', 527 
Selenic acid, 178 

Selenious anhydride and acid, 177 
Selenite, 394 
Selenium, 174 

— chlorides, 179 

— compounds, 175 

— •disulphide, 178 
Senarmontite, 253 
Serpentine, 271, 413, 417 
Sesqui-oxides, 55 
Siderite, 436, 446 
Siemens-Martin process, 441 
Silica, 268 

Siliceous calamine, 493 

— sinter, 269 
Silicic acid, 271 
Silico-chloroform, 276 
Silico-formic anhydride, 277 
Silicon, 266 

— compounds, 268 

— fluoride, 274 

— hexachloride, 274 

— hexiodide, 274 

— hydride, 272 

— tetrabromide, 274 

— tetrachloride, 273 

— tetriodide, 274 
Silicotungstic acid, 328 
Silicoxalic acid, 277 
Silver, 543 

— alloys of, 551 

— arsenate, 547 

— arsenite, 547 

— bromide, 548 

— carbonate, 547 

— chloride, 548 

— chromate, 548 

— compounds, 545 

— copper-glance, 543 

— cyanide, 550 


604 


Index. 


i. 


SIL 


SUL 


Silver, detection of, 551 

— fluoride, 550 
Silver-glance, 543, 551 
Silver iodide, 548 

— nitrate, 546 

— oxide, 545 

— peroxide, 546 

1— phosphate, 547 

— sulphate, 546 

— sulphide, 550 
Skutterudite, 489 
Slaked lime, 390 
Smalt, 489 
Smaltite, 229, 484 
Smithsonite, 493 
Soapstone, 413 
Soda, 369 

— felspar, 425 

— sesquicarbonate of, 374 
Sodium, 365 

— acid carbonate, 374 

— acid sulphate, 368 

— acid sulphite, 368 

— aluminate, 422 

— arsenate, 375 

— auric chloride, 553 

— borate, 376 

— carbonate, 369 

— chloride, 376 

— chlorplatinate, 562 

— chromate, 477 

— detection of, 378 

— diacid phosphate, 375 

— ferrocyanide, 456 

— hydrate, 367 

— metaphosphate, 375 

— metastannate, 519 

— monacid phosphate, 374 

— nitrate, 369 

— nitroprusside, 458 

— oxide, 367 
—■ peroxide, 367 

— phosphates, 375 

— pyrophosphate, 375 

— stannate, 519 

— sulphate, 367 

— sulphite, 368 

— thiosulphate, 368 
Solder, 521 
Solution, 76 


Solution, supersaturated, 84 
Sombrerite, 396 
Spathic iron ore, 286, 436, 446 
Spectrum analysis, 337 
Specular iron, 436, 447 
Speculum metal, 531 
Speiss-cobalt, 484, 489 
Sphene, 323 
Spiegeleisen, 439 
Spinelle, 422 
Spongy platinum, 558 
Stahl’s phlogiston theory, 31 
Stannic acid, 520 

— chloride, 520 

— compounds, 519 

— hydrate, 519 

— oxide, 519 

— sulphide, 520 
Stannous chloride, 518 

— compounds, 517 

— oxide and hydrate, 517 

— sulphate, 517 

— sulphide, 518 
Steatite, 413, 417 
Steel, 440 
Stibnite, 258 
Stoneware, 426 
Strontianite, 286, 406 
Strontium, 406 

— carbonate, 407 

— chloride, 407 

— chromate, 477 

— detection of, 408 

— hydrate, 407 

— nitrate, 407 

— oxide, 406 

— peroxide, 407 

— sulphate, 407 

— sulphide, 407 
Struvite, 416 
Suboxides, 55 
Sugar of lead, 505 
Sulphantimonic acid, 259 
Sulphites, 154 

Sulpho-acids, -bases, and -salts, 57 
Sulphocarbonic acid, 302 
Sulphocyanic acid, 317 
Sulphur, 139 

— compounds of, with oxygen, 148 

— dichloride, 173 


Index. 


605; 


SUL 

Sulphur diselenide, 178 

— flowers of, 142 

— group of elements, 137 

— hexiodide, 174 

— milk of, 145 

— tetrachloride, 174 
Sulphuric acid, 155 

— fuming or Nordhausen, 163 
Sulphuretted hydrogen, 145 
Sulphurous acid, 154 

— anhydride, 150 
Sulphuryl, 152 

— chloride, 164 

Superphosphate of lime, 209, 398 
Supersaturated solutions, 84 
Sylvanite, 552 


Tables, 586 
Talc, 413, 417 
Tantalite, 330 
Tantalum, 330 
Tartar emetic, 254 
Tellurium, 179 

Temperature, influence of, on che¬ 
mical affinity, 52 
Terbium, 433 
Terne-plate, 443 
Tetrads, 47 
Tetradymite, 179 
Tetrahedrite, 523 
Tetrathionic acid, 172 
Tetravalent elements, 47 
Thallic compounds, 512 
Thallium, 509 
—- compounds, 511 

— detection of, 512 
Thallous compounds, 511 
Theory, atomic, 44 

— electro-chemical, 46 

— phlogiston, 31 
Thiosulphuric acid, 167 
'1 horium, 434 

Tin, 516 

— alloys of, 521 

— compounds, 517 

— detection of, 522 
Tin-foil, 521 
Tin-plate, 443, 521 
Tin-salt, 518 


VAN 

Tin-stone, 516, 519 

Tinkal, 260, 376 

Titanic anhydride and acid, 222 

— iron, 323 
Titanite, 323 
Titanium, 322 

— nitride, 325 

— sesquioxide, 324 

— tetrachloride, 324 

— tetrafluoride, 325 
Tombac, 531 
Triads, 47 
Tribasic acids, 56 
Tridymite, 269 
Triferric tetroxide, 449 
Trimanganic tetroxide, 464 
Triphylline, 379 
Trithionic acid, 172 

Trivalent elements, 47 
Trona, 286, 374 
Tungsten, 327 

— chlorides, 329 

— steel, 329 

Tungstic anhydride and acid, 328* 
Turnbull’s blue, 458 
Turpeth mineral, 535 
Type metal, 252, 507 


Ultramarine, 427 
Uranates, 482 
Uranic carbonate, 481 

— nitrate, 480 

— oxide, 480 

— phosphate, 481 

— sulphate, 480 
Uranium, 479 
Uranoso-uranic oxide, 482. 
Uranous chloride, 482 

— — oxide, 482 

— sulphate, 482 
Uranyl, 481 

— chloride, 481 

— sulphide, 482 


Valency, 47 
Valentinite, 253 
Vanadinite, 329 
Vanadium, 329 


6 o6 


Index. 


VER 

Verdigris, 524, 528 
Vermilion, 539 
Vivianite, 447 
Volumes, law of, 43 


Wad, 460 
' Water, 80 

— hard, 83 

— of constitution, 86 

— of crystallization, 85 
Waters, mineral, 82 
Wavellite, 425. 
Weights, atomic, 60 

— equivalent, 42 
White arsenic, 230, 233 

— lead, 504 

— precipitate, 537 

— vitriol, 495 
Willemite, 493 
Witherite, 286, 408 
Wolfram, 327 
Woulffs bottle, 19 
Wrought iron, 440 
Wulfenite, 325 


ZIR 

Yellow prussiate, 453 
Ytterbium, 433 
Yttrium, 433 
Yttrotantalite, 433 


Zaffre, 489 
Zeolites, 425 
Zinc, 493 

— blende, 493, 49 ° 

— carbonate, 495 

— chloride, 495 

— detection of, 496 

— hydrate, 495 

— nitrate, 495 

— oxide, 494 

— spar, 495 

— sulphate, 495 

— sulphide, 496 

— white, 495 
Zinkenite, 507 
Zircon, 434 
Zirconium, 434 

— compounds, 435 


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